UC-NRLF 


SB    35    M3E 


LABORATORY  DIRECTIONS 

FOR 

ELEMENTARv  CHEMISTRY 

MATT  ILL 


jL^       J 


LABORATORY   DIRECTIONS 


FOR 


ELEMENTARY  CHEMISTRY 


DESIGNED  TO  ACCOMPANY 

A  TEXTBOOK  OF  CHEMISTRY 
BY  W.  A.  NOYES 


BY 
HELEN  ISHAM  MATTILL,  PH.  D. 

FORMERLY   ASSOCIATE   IN   CHEMISTRY     AT  THE   UNIVERSITY   OP   ILLINOIS 


NEW  YORK 

HENRY  HOLT  AND  COMPANY 
1914 


COPYRIGHT,  1912,  1914 

BY 
HENRY  HOLT  AND  COMPANY 

IN  MEMORIAM 


THE. MAPLE- PRESS. YORK. PA 


PREFACE 

The  importance  of  the  laboratory  work  in  a  course  in  elemen- 
tary chemistry  cannot  be  too  strongly  emphasized.  It  is  only 
by  the  actual  contact  and  experimentation  with  some  of  the 
many  materials  described  in  a  text-book  that  the  student  ever 
comes  to  any  realizing  sense  of  the  means  by  which  the  science  of 
chemistry  has  been  built  up.  For  this  reason  it  is  advisable 
to  supplement  the  laboratory  exercises  with  discussions  of  an 
explanatory  nature,  in  which  the  correlation  between  laboratory 
and  text-book  work  may  be  brought  out. 

The  experiments  described  in  this  guide  have  been  chosen 
with  the  following  objects  in  mind:  To  reduce  the  variety  of 
materials  handled,  consequently  the  confusion  of  many  new 
names  and  strange  materials,  to  a  minimum;  to  make  each 
experiment  a  fairly  exhaustive  study  of  some  particular 
material  or  property;  to  have  the  sequence  of  experiments 
such  as  will  lead  to  a  certain  amount  of  reasoning  from  analogy. 
This  guide  does  not  pretend  to  be  an  exhaustive  description  of 
elementary  experiments,  but  rather  a  selection  of  a  few  suitable, 
consecutive  experiments  from  the  many  possible. 

I  wish  to  take  this  opportunity  to  acknowledge  the  many 
helpful  suggestions  and  criticisms  which  I  have  received  from 
Drs.  C.  W.  Balke,  C.  H.  Hecker,  S.  B.  Hopkins  and  W.  A.  Noyes 
in  the  preparation  of  this  book. 


in 


CONTENTS 

PAGH 

GENERAL  DIRECTIONS 1 

CHAPTER 

I.  APPARATUS  AND  MANIPULATION 10 

II.  PURE  SUBSTANCES  AND  MIXTURES 16 

III.  OXYGEN 21 

IV.  HYDROGEN 26 

V.  WATER  AND  HYDROGEN  PEROXIDE 31 

VI.  THE  HALOGEN  FAMILY 38 

VII.  SULFUR 55 

VIII.  NITROGEN 63 

IX.  THE  ATMOSPHERE 68 

X.  PHOSPHORUS,  ARSENIC,  ANTIMONY  AND  BISMUTH 71 

XI.  CARBON 76 

XII.  SILICON 81 

XIII.  BORON 83 

XIV.  METALS  AND  NON-METALS 85 

XV.  THE  ALKALI  METALS 89 

XVI.  COPPER  AND  SILVER 92 

XVII.  CALCIUM,  STRONTIUM  AND  BARIUM 95 

XVIII.  MAGNESIUM,  ZINC  AND  MERCURY 98 

XIX.  ALUMINIUM      102 

XX.  LEAD  AND  TIN 104 

XXI.  CHROMIUM 107 

XXII.  MANGANESE 109 

XXIII.  IRON Ill 

INDEX                                                                                                             .  115 


LABORATORY  DIRECTIONS 

FOR  ELEMENTARY 

CHEMISTRY 

GENERAL  DIRECTIONS 

Work  in  the  laboratory  should  be  undertaken  with  the  fol- 
lowing objects  in  mind:  1.  By  direct  handling  of  the  materials 
to  become  familiar  with  certain  "chemical  reactrons.  2.  By 
analogy,  and  text-book  study,  to  relate,  the  experiments 
actually  performed  in  the  i3br3raiior^>>withjr'as>'maiiy  facts 
brought  up  in  the  lecture  and  quiz  room  discussion  as  possible. 
3.  By  handling  and  manipulating  apparatus  to  become  familiar 
with  the  tools  at  the  disposal  of  the  chemist,  and  the  proper 
use  of  the  same,  and  at  the  same  time  to  acquire  ability  and 
dexterity  in  their  use.  A  student,  however  good  his  under- 
standing of  chemical  facts,  has  not  reached  the  goal  unless 
he  is  able  to  perform  neatly  and  properly  any  given  chemical 
manipulation.  4.  To  acquire  habits  of  observation  and  a 
proper  discrimination  between  important  and  unimportant 
details,  and  to  draw  conclusions  of  a  general  nature  from  a 
specific  case.  5.  To  record  accurately,  briefly,  and  promptly 
the  phenomena  observed. 

In  the  pursuit  of  these  objects  the  keeping  of  a  laboratory 
note-book  is  of  prime  importance.  It  is  absolutely  essential 
that  these  notes  be  recorded  at  the  time  the  work  is  performed. 
They  should  give  a  clear  but  brief  account  of  work  done  and 
should  include  a  statement  of  the  object  of  the  experiment,  a 
record  of  all  phenomena  observed,  such  as  change  of  color, 
appearance  of  a  precipitate  or  a  gas,  etc.,  and  answer  all  ques- 

1 


2  GENERAL  DIRECTIONS 

tions.  A  drawing  of  the  apparatus  used  may  often  reduce 
the  length  of  the  description  required.  Wherever  possible  a 
chemical  equation  should  be  written  to  express  the  change 
taking  place.  In  order  to  write  such  equations  it  may  be  nec- 
essary to  refer  to  the  text-book  for  the  products  of  the  reaction. 
After  the  products  have  been  determined  the  equation 
should  be  written  and  balanced  without  reference  to  text- 
book or  other  aid.  Do  not  delay  in  recording  the  observations 
made  during  a  given  experiment;  each  change  should  be  jotted 
down  as  observed.  It  is  more  essential  that  the  note-book  be 
a  record  of  the  experiment  taken  at  the  time  it  is  performed 
than  that  it  should  be  neat  and  well  written,  though  both  of 
these  ends  may  be  accomplished  at  once  without  too  great 
difficulty.  Under  no  circumstances  should  the  record  of  an 
experiment'  jb£,  placed  'Qi)  loose  paper,  the  loss  of  which  would 
mean  the  repetition  of  the  experiment  so  recorded. 

The  f pJlcwfiig, -gejner&l  directions  for  manipulation  should  be 
closely  followed: 

(1)  Do  not  use  too  much  material.     The  amount  called  for 
in  any  experiment  is  ample  to  secure  the  results  desired.     Tak- 
ing larger  amounts  of  material  is  wasteful,  frequently  dangerous, 
and  always  results  in  loss  of  time. 

(2)  Keep  all  bottles  in  their  proper  places. 

(3)  Material  spilled  on  the  side  shelf  should  be  promptly 
swept  into  a  waste  jar. 

(4)  Never  place  the  stopper  of  a  reagent  bottle  on  the  table — 
hold  it  between  the  ringers. 

(5)  Do  not  dip  stirring  rods  or  tubes  into  the  reagent  bottles. 
Pour  a  little  of  the  reagent  into  a  test  tube. 

(6)  Do  not  waste  gas  or  water. 

(7)  In  pushing  a  glass  tube  through  a  stopper  wet  both  the 
tube  and  the  stopper.     It  is  wise  also  to  protect  the  hand  by 
covering  the  tube  with  a  towel. 

(8)  Do  not  heat  beakers  or  flasks  with  the  direct  flame,  but 
always  place  them  on  a  wire  gauze.     Test  tubes  are  the  only 
glass  receptacles  which  can  be  heated  in  the  direct. flame. 


GENERAL  DIRECTIONS  3 

(9)  Before  leaving  at  the  end  of  a  laboratory  period  see  that 
all  apparatus  is  in  its  proper  place,  and  that  the  desk  is  clean. 

DIRECTIONS  FOR  QUANTITATIVE  EXPERIMENTS 

In  performing  quantitative  experiments  the  student  must 
keep  in  mind  the  object  of  the  given  experiment,  and  the  fact 
that  successful  results  in  quantitative  work  depend  largely 
on  the  care  and  accuracy  with  which  given  manipulations  are 
carried  out.  A  given  experiment  may  be  successful  in  the 
hands  of  one  student  and  a  complete  failure  in  the  hands  of 
another,  not  because  of  any  fault  in  the  choice  of  the  experi- 
ment, but  because  of  the  differences  in  manipulation  which 
result  from  an  incomplete  understanding  of  the  object  of  the 
experiment,  and  of  the  possibilities  of  error  and  how  to  avoid 
them.  For  this  reason  it  is  especially  essential  that  the  direc- 
tions for  an  experiment  marked  quantitative  should  be  studied 
from  beginning  to  end  before  any  manipulation  is  started.  A 
student  should  begin  the  manipulation  of  a  quantitative 
experiment  with  the  object  of  the  experiment  so  clearly  in 
mind  that  reference  to  the  directions  is  unnecessary  except  as 
regards  the  details  of  manipulation. 

The  record  of  a  quantitative  experiment  should  be  kept  in 
tabulated  form,  and  should  in  every  case  be  kept  in  the  note- 
book. If  the  experiment  has  to  be  repeated  the  record  of  every 
succeeding  trial  must  also  be  in  the  note-book.  On  reading 
over  the  directions  make  a  note  of  what  measurements  are  to 
be  taken  and  in  what  order  they  will  be  used,  then  prepare  a 
sheet  of  the  laboratory  note-book  with  the  proper  headings  in 
the  proper  order  before  starting  the  experiment.  For  example, 
in  the  first  experiment  (No.  8)  the  weights  of  two  tubes  empty, 
with  potassium  chlorate,  and  with  the  residue  on  heating  potas- 
sium chlorate  are  to  be  determined  and  from  that  data  the 
percentage  of  potassium  chloride  obtained  on  heating  a  given 
amount  of  potassium  chlorate  is  to  be  calculated.  The  headings 
then  might  be  arranged  as  follows: 


GENERAL  DIRECTIONS 

Small  tube      Large  tube 


Tube  +  potassium  chlorate, 
Tube  empty, 
Potassium  chlorate, 

Tube  +  potassium  chloride  (residue), 
Tube  empty, 
Potassium  chloride, 

Per  cent  of  potassium  chloride  in  potassium 
chlorate, 


The  arrangement  of  the  data  sheet  before  performing 
the  experiment  is  a  good  test  of  the  student's  understanding  of 
the  object  of  the  experiment,  and  for  that  reason  should  be 
recorded  in  the  note-book  before  the  manipulation  is  begun. 

While  the  directions  for  each  quantitative  experiment  de- 
scribe the  special  precautions  to  be  observed  in  the  given  experi- 
ment there  are  other  precautions  which  apply  to  all  quantitative 
work  and  which  should  be  kept  constantly  in  mind,  of  which 
the  following  are  the  most  important. 

Weighing. — In  all  quantitative  experiments  use  the  chemical 
balance.  To  prepare  a  dish  for  weighing  it  should  be  carefully 
cleaned  and  wiped  dry  with  a  clean  towel.  After  a  dish  has 
been  prepared  for  weighing  it  should  be  handled  as  little  as 
possible  with  the  fingers  because  of  deposition  of  moisture  and 
oil  from  the  skin.  Material  may  be  placed  in  a  dish  so  pre- 
pared arid  the  required  processes  carried  out  according  to  the 
directions  of  the  given  experiment,  but  the  student  should 
always  keep  in  mind  the  fact  that  an  error  may  be  introduced 
into  the  weight  of  any  given  material  by  allowing  foreign  matter 
to  fall  by  accident  into  the  dish,  or  by  accidental  loss  of  material 
from  the  dish.  Directions  to  avoid  loss  of  material  by  spatter- 
ing, etc.,  must  be  carefully  observed.  The  success  of  every 
quantitative  experiment  depends  on  determining  by  weighing, 
measuring  or  otherwise,  all  of  a  given  material  which  results 
from  a  given  chemical  process. 

Testing  for  Leaks. — Some  of  the  quantitative  experiments 
involve  the  evolution  and  measurement  of  a  volume  of  gas. 


GENERAL  DIRECTIONS  5 

In  all  such  experiments  the  apparatus  must  be  constructed 
with  unusual  care  and  tested  to  determine  that  there  is  no 
possibility  of  a  loss  of  gas  through  leaks.  Rubber,  instead  of 
cork  stoppers  should  be  used  and  wherever  rubber  tubing  is 
used  in  making  connections  the  glass  tubes,  which  should  always 
be  fire  polished  (see  Exp.  3  a),  should  meet  inside  the  rubber 
tubing.  Unless  an  apparatus  is  air-tight  it  should  not  be  used  in  a 
quantitative  experiment. 

STANDARD  CONDITIONS  FOR  GASES 

In  order  that  gas  volumes  may  be  comparable  they  must  be 
given  under  fixed  conditions,  since  the  volume  of  a  given  sample 
of  gas  varies  with  both  the  temperature  and  the  pressure  under 
which  it  exists.  The  fixed  conditions  which  have  been  arbi- 
trarily chosen  by  chemists  are  that  the  gas  should  be  dry, 
under  the  pressure  of  760  mm.  of  mercury  measured  at  0°  C., 
and  at  a  temperature  of  0°  C.  The  gas  as  obtained  in  the 
laboratory  is  usually  saturated  with  moisture,  though  in  certain 
cases,  for  example  Exp.  20,  it  may  be  measured  dry,  and  is 
under  the  temperature  and  pressure  conditions  obtaining  in  the 
laboratory.  These  conditions  may  be  determined  by  a  ther- 
mometer and  barometer.  The  barometer  reading  may  indicate 
a  certain  atmospheric  pressure,  but  it  is  self-evident  that  if  the 
temperature  of  the  room  were  lower  the  same  atmospheric 
pressure  would  give  a  lower  barometric  reading,  because  of  the 
contraction  of  the  mercury  on  cooling.  For  this  reason  the 
barometric  reading  must  be  corrected  for  the  temperature  at 
which  the  barometer  stands.  A  table  of  corrections  for  ordi- 
nary temperatures  can  be  readily  prepared  from  the  coefficient 
of  expansion  of  mercury,  and  may  be  found  below.  To  correct 
a  barometric  reading  to  0°  C.,  then,  the  observed  reading  should 
be  reduced  by  the  amount  indicated  in  the  table,  for  the 
observed  temperature.  For  instance  suppose  the  barometer 
reads  74.3  cm.  at  17°  C.,  the  correction  for  17°  C.  is  2.2  mm., 
therefore  the  barometric  reading  at  0°  C.  would  be  743  —  2.2  = 


6  GENERAL  DIRECTIONS 

740.8  mm.  But  if  the  gas  was  saturated  with  moisture  the 
pressure  indicated  by  the  barometer  was  being  divided  between 
the  two  components  of  the  mixture,  the  water  vapor  and  the 
gas.  Since  the  pressure  of  water  vapor  at  varying  temperatures 
has  been  accurately  determined,  that  may  be  obtained  from  a 
table,  and  the  pressure  of  the  gas  in  the  mixture  is  then  the 
difference  between  the  barometric  reading  and  the  pressure 
of  water  vapor  at  the  observed  temperature.  For  instance,  the 
corrected  barometric  reading,  as  found  above,  was  740.8  mm., 
but  the  gas  in  question  was  saturated  with  moisture.  The 
vapor  pressure  of  water  at  17°  C.  is  found  from  the  table  to  be 
14.5  mm.,  therefore  the  pressure  of  the  dry  gas  was  740.8  —  14.5 
=  726.3  mm.  of  mercury  measured  at  0°  C.  The  gas  volume 
must  then  be  reduced  to  standard  conditions  of  temperature 
and  pressure  by  applying  the  laws  of  Boyle  and  Charles,  before 
any  comparison  of  volume  may  be  made. 

TABLE  I 

TEMPERATURE  CORRECTIONS  FOR  REDUCING  A  COLUMN  OF  MERCURY  750  MM. 
HIGH  WHEN  READ  ON  A  GLASS  SCALE,  TO  0°  C.,  IN  MILLIMETERS 

Temperature,        15       16       17       18       19       20       21       22      23      24       25 
Correction,  2         2.1     2.2      2.3     2.5      2.6     2.7     2.8    3.0     3.1      3.2 

Temperature,        26       27       28      29      30       31       32       33     34      35 
Correction,  3.4     3.5     3.6     3.7     3.9     4.0     4.1      4.3    4.4     4.5 

TABLE  II 
VAPOR  PRESSURE  OF  WATER  IN  MM.  OF  MERCURY 


Temperature, 

0° 

5 

10 

15 

16 

17 

18 

19 

Pressure, 

4.6 

6, 

5 

9. 

2 

12. 

8 

13.6 

14.5 

15.5 

16.5 

Temperatures, 

20 

21 

22 

23 

24 

25 

26 

27 

Pressure, 

17.5 

18. 

.7 

19 

,8 

21. 

1 

22.4 

23.8 

25.2 

26.7 

Temperature, 

28  • 

29 

30 

31 

32 

33 

34 

35 

Pressure, 

28.4 

30, 

1 

31, 

,8 

33. 

7 

35.7 

37.7 

39.9 

42.2 

Temperature, 

40 

45 

50 

55 

60 

65 

70 

75 

Pressure, 

55.3 

71. 

9 

92. 

5 

118. 

1 

149.5 

187.6 

233.8 

289.3 

Temperature, 

80 

85 

90 

95 

100 

Pressure, 

355.5 

433. 

,8 

526 

.0 

634. 

0 

760 

GENERAL  DIRECTIONS 


TABLE  III 

DENSITIES  AND  PER  CENT  COMPOSITION  OF  WATER  SOLUTIONS  OF  ACIDS 
AND  AMMONIA.     (Taken  from  Landolt,  Bornstein  and  Roth) 


Acetic 

acid 

Nitric  acid                            Phosphoric  acid 

Sp.  Gr. 

%  acid 

Sp.  Gr. 

%  acid                        Sp.  Gr. 

%acid 

=  d20 

I15 

d17'5 

4 

4 

4 

1.012 

10 

1.06 

10.67                           1.06 

10 

1.026 

20 

1.12 

20.22                            1.12 

20 

1.038 

30 

1.19 

30.87                            1.18 

30 

1.049 

40 

1.25 

39.8 

.26 

40 

1.057 

50 

1.31 

49.05 

.34 

50 

1.064 

60 

1.37 

59.36 

.43 

60 

1.069 

70 

1.39 

63.2 

.53 

70 

1.070 

80 

1.42 

69.77 

.65 

80 

1.066 

90 

.76 

90 

1.050 

100 

Maximum  density  is  at 
78%  acid. 


Hydrochloric  acid 

Sulfuric  acid 

Sp.  Gr. 

%acid 

Sp.  Gr. 

%acid 

^4 

4 

1.05 

10.17 

1.068 

10 

1.10 

20.00 

1.14 

20 

1.15 

29.57 

1.22 

30 

1.16 

31.52 

1.30 

40 

1.20 

39.11 

1.40 

50 

1.50 

60 

1.61 

70 

1.73 

80 

1.82 

90 

1.84 

100 

Ammonia 
Sp.  Gr.      %NH3 


,15 
'd"4 
0.96 
0.92 
0.90 


10 
20 

28.5 


8  GENERAL  DIRECTIONS 

TABLE  IV 

SOLUBILITIES.  (Taken  from  Landolt,  Bornstein  and  Roth.)  Giving 
in  order  the  formula  of  the  salt  or  its  hydrate  stable  at  10°,  the  weight  in 
grams  of  the  salt,  or  the  hydrate  (always  calculated  from  the  formula  given 
in  column  1)  required  to  make  a  liter  of  normal  solution,  and  the  weight  of 
anhydrous  substance  in  100  grams  of  a  saturated  solution  at  10°,  20°,  30°, 
40°,  50°  and  100°. 


Formula 

Grams  per 
liter  for  a 
normal 
solution1 

Solubility 

10° 

20°               30° 

40° 

50° 

100° 

AgC2H3O2  
AffTl 

166.9 
143.3 
187.8 
234.8 
169.9 
155.9 
241.45 
111.1 
118.7 
98.7 
140.2 
157.7 
50. 
109.5 
37. 
86. 
114.1 
154.3 
119. 
145.6 
50. 
147  9 

0.875 
8.9-10-5 

1.037 
1.53-10-4 
0.84-10-5 
3  53-10"7 

1.215 

1.413 

1.637 

2.19-10-8 
3.7-10-4 

AeBr 

Agl 

AgNOs     

61.5 

68.3 
0.78 

73.            |77.             80. 

30.1 
1.46 

Ag2SO4  
A1C13-6H2O             ! 

A12(S04)3-18H26.. 
A12K2(SO4)4-24H2O 
BaCOs 

25.1 

26.6 

28.8 

7.74 

31.4 

34.3 

47.1 
30.6xaq. 

1.8-10-3 
25. 
2.17 

. 

BaCl2-2H2O 

26.3              ,27.6 
3.36               4.75 

29.- 
6.85 

30.4          '37. 
10.5           
2-10-s 

BaO9H2O  
CaCOs              .  .  .  . 

CaCl2-6H2O  
CaOH2O 

39.4 

42.7 
0.123 
0.20 

53.54aq 
0.113        0.104 
0.21           0.21 
56.3          57.51aq 
58.4          ;61.4 
36.1          39.4 

61.42aq 
0.096        0.05 

CaS(V2H26  
CdCl2-2iH2O  .... 
Cd(NO3)2-4H2O.... 
CoCl2-6H2O  
Co(NOs)2-6H2O.... 
CrOs  
Cu(NOs)2'6H2O 

0.192 
47.4 

is'.s"" 

59.5 
o6'.7'2aq 

31. 

33.3 

\fi9,  h 

64.6 

67.4 

55.6 
47.9 

51.6" 

73.25aq 
24.8 
6.4 

61.53aq 

FeCl3-6H2O  { 

FeSO4'7H2O  
H3B03  
HgCl 

90.              i45. 

75.9 

28.7 
8. 
7-10~4 
8.7 
43  2 

75.9 
32.7 
10.3 

84.  2  an" 

139. 
21. 
236. 
135.6 
119. 
167. 
87. 
74.6 
122.6 
138.6 
97  .2(64.8)  4 
147.  (49)  4 
50.2 
166. 
214. 
158.  (31.6)  4 
101. 
87.2 
101.7 
128.3 
123.3 
99. 
120.6 
53.5 
39.62 
80.1 
66.1 

191.2 

139. 
151. 

17. 

21. 
4.9 
0.38  -10-* 
6.9 
39.4 
6.54 

28.2 

HgCl2 

6.19 
38. 



10.2 

35. 
51.2 
33.2 
60.9 
36. 
35.9 
15.8 
44.2 
50.5 

KBr 

KBrOs 

11.7 

53.9 

28.7 
12.7 

54!s" 
30. 
16.5 
5. 
40.8 

K2CO3-2H2O 

53.3 

27.2 
9.2 

KC1 

23.8 

4.7 

25.5 

6.8 
1.9  (25°) 

KC1O3 

KC1O4 

K2CrO4.  .  . 

37.9 

7.8 
21.7 
57.7 

38.6 
11.6 
24.9 
59.1 
7.5 
6. 
24.1 
10. 
35.3 
42. 
26.2 

39.5 
15.4 
28.1 
60.4 
10.5 
8.3 
31.6 
11.5 

40.1 
22.6 
31.2 
61.5 
11.4 
10.4 
39.2 
13. 
3fi   5 

K2Cr2Or 

KHCO3 

34.2 
62.7 
24.4 
14.3 
46.3 
14. 

KI... 

67.6 

KIOs 

KMnO4 

4.01 
17.7 
8.4 
34.9 

KNOs.  .  . 

71.1 
19.4 
42.2 

• 

K2SO4 

MgCl2-6H2O  
Mg(NOs)2-6H2O. 
MgS04'7H20.... 
MnCl2-4H2O  
MnSO45H2O.... 
NH4C1 

45  .  9 

23.6 

29. 

44  7 

31. 

33.56aq 
49.5 
37.3 
33.5 

42.5 
53  .  7  2aq 
24.9 
43.6 

38.6 
27.1 

17.4 

43'." 

39.4  laq 
29.3 
21.3 
70.8 
43.8 
3  7 

25. 
13.7 

42i2" 
1.58 

31.4 

NH4HCO3 

NH4N03  

(NH4)2SO4 

74.8 
144.8 

78.5 
45.8 
9  5 

89.7 
50.8 
35.  5aq 
55.  an 
47.6 

Na2B407-10H20.  . 
NaBr-2H2O  
NaBrOs  

47.5 

27.7 

51.4 
33.4 

53.7 

GENERAL  DIRECTIONS 

TABLE  IV. — Continued 


Formula 

Grams  per 
liter  for  a 
normal 
solution 

Solubility 

10° 

20° 

30° 

40° 

50° 

100° 

Na2CO3.10H2O... 
NaCl  

NaClOs 

143. 
58.5 
106.5 
171.2 
149.1 
42.2 
85. 
126. 
161.2 
124. 
145.5 
138.9 
166. 
73.8 
133.3 
141.8 
132.8 
91.8 
91.1 
148.8 
143.8 

11.2 
26.3 

17.6 
26.4 
49.7 
44. 
64. 
8.8 
46.8 
20.5 
16. 
41. 
49. 
0.96 
34. 
1-10'8 
35. 
41.5 
0.7 
15-10-3 
78.6  IJaq 
54. 

29. 
26.5 

47.  4aq 
66.3 
10. 
49. 

33.2  laq 
26.7 
56.5 
49. 
68.8 
11.3 
51. 

32.2 
26.8 

31. 
28.1 
67.1 
55  .  8  an 
82.  an 

Na2CrOrlOH20... 
Na2Cr207'2H2O  .  . 
NaHCOs 

33.4 
63. 
7.6 
44.6 
16.6 
8.3 
38. 

51. 
71.3 
12.7 
53. 

NaNO3  

64.4 
24  .  8  an 
29.9 
72.7 

Na2SO3'7H2O  
Na2S04'10H2O... 
Na2S2O.r5H2O  ... 
Ni(\03)2-6H2O..  . 
PbCl2 

29. 

46. 

32.5  an 
50.6 
55. 

31.8 
63.  2aq 

3.2 
56. 

Pb(NO3)2 

30.8 

38. 

41. 

44. 

SrCOs  
SrCl2-6H2O  

32.6 
35.5 
0.48 
9.9-10-3 
73.1 

37  .  5 
46.7 
1. 

40. 

47.7  an 
1.5 

42.7 
48.1 
2.1 
16-  10-3 

50.5 
50.3 
19.5 
18-10-3 

86. 

44  !  'laq  ' 

Sr(N03)2-4H20.... 
Sr(>9H2O  

SrSO4  

ZnCl2-2  1/2  H2O  . 
Zn(NO3)2-6H2O.... 
ZnSO4-7H2O  

82    an 

67.4  3aq 
41.2 

43'.5'6aq 

32.3 

1  In  many  cases  this  quantity  exceeds  the  solubility  of  the  salt,  but  it  is  always  given 
because  of  the  chemical  significance  of  the  value. 

2  This  is  half  the  formula  weight. 

3  Anhydrous. 

4  The  figure  in  parenthesis  applies  when  the  salt  is  to  be  used  as  an  oxidizing  agent  in 
acid  solution. 

TABLE  V 

SOLUBILITY  AND  DENSITY  OF  GASES.  (Taken  from  Landolt,  Bornstein 
and  Roth.)  Giving  the  volume  of  gas,  measured  at  standard  conditions, 
absorbed  by  one  volume  of  water,  at  0°,  20°,  50°  and  100°  and  the  weight 
per  liter  of  the  gas  under  standard  conditions. 


Gas 

H2 

N2 

02 

CL 

HC1 

Br2 

HBr 

NH3 

NO 

N2O 

CO  2 

S02 

H2S 

CH4 

C2H4 

C2H2 


0° 


20° 


50° 


100° 


0.0215   0.0184 
0.0239   0.0164 
0.0489   0.0310 
4.61     2.26 
506.9    442  3 

0.0161 
0.0106 
0.0209 
1.2 
361 

0.0160 
0.0100 
0.0170 
0.000 

60.5     21.3 
612.5      

6.5 

468 



1305     715.4 

0.0738   0.047 
0.629 

0.0315 

0.0262 

1.713    0.878 
79  8     39  4 

0.436 



4.6     2.5 
0.0556   0.033 
0.226    0.122 

1.4 
0.0213 

0.8 
0.017 

1  .  73     1  .  03 

Weight 
per  liter 
0.0898 
1.256 
1.429 
3.180 
1.610 
7.143 
3.616 
0.762 
1.340 
1.970 
1.977 
2.870 
1.542 
0.717 
1.252 
1.162 


CHAPTER  I 
APPARATUS  AND  MANIPULATION 

Check  up  the  apparatus  found  in  the  desk,  using  a  list 
obtained  from  the  storeroom.  Record  any  missing  articles  on 
a  sheet  furnished  for  that  purpose,  and  after  getting  the  approval 
of  an  instructor  take  both  lists  to  the  storeroom  and  obtain  the 
missing  articles. 

Study  the  apparatus  supplied  with  a  view  to  determining  its 
uses  and  proper  manipulation  as  follows: 

1.  Bunsen  Burner. 

a.  Examine  the  construction  of  the  Bunsen  burner.     Explain 
the  function  of  each  part. 

b.  Attach  the  burner  to  the  gas  supply  by  means  of  a  rubber 
tube,  open  the  holes  at  the  bottom  of  the  vertical  tube  and  light 
the  gas.     What  is  the   character  of  the  flame?     Show  the 
structure  of  the  flame  by  a  sketch.     Explore  the  various  parts 
of  the  flame  with  a  platinum  wire.     What  are  the  relative 
temperatures  of  the  different  parts?     Where  should  an  object 
be  placed  to  secure  the  maximum  heating  effect?     Could   a 
flame  of  this  character  be  used  for  illuminating  purposes? 
What  is  the  principle  of  the  Welsbach  light? 

c.  Close  the  holes  at  the  bottom  of  the  tube.     What  is  the 
character  of  the  flame  now?     Hold  a  cold  porcelain  dish  in  the 
flame  for  a  moment.    What  is  the  deposit?     (Illuminating  gas 
contains  a  compound  of  carbon  and  hydrogen.)     Why  is  it 
possible  to  obtain  a  deposit  of  this  substance  on  a  cold  dish? 
Hold  a  clean  portion  of  the  dish  above,  but  not  in,  the  flame. 
Why  is  no  deposit  obtained  in  this  case  ?     What  causes  the 
luminosity  of  the  flame? 

10 


BUNSEN  BURNER 


11 


d.  Repeat  c   using  a  non-luminous    flame.      Explain  the 
result. 

e.  Extinguish  the  flame,  turn  on  the  gas  without  lighting  it, 
and  hold  the  dish  in  the  current  of  gas.     Why  is  no  deposit 
obtained? 

/.  Light  the  gas  and  show  the  structure  of  the  luminous  flame 
by  a  sketch  (turn  the  flame  low  to  avoid  flickering).  Test  the 
temperature  of  this  flame  with  a  platinum  wire.  How  does  it 
compare  with  the  temperature  of  the  non-luminous  flame  ? 

g.  Quickly  introduce  the  head  of  a  match  into  the  innermost 
cone  of  both  flames.  What  does  this  experiment  indicate  as  to 


Burrette 


7 


Pipette 
FIG.  1. 


FIG.  1. 

the  temperature  of  the  central  cone?  What  composes  this 
central  portion?  Devise  and  perform  an  experiment  to  test 
this  point. 

2.  Measuring  Instruments. 

The  measuring  instruments  consist  of  the  cylinder,  pipette 
and  burette  (side  shelf).  Fig.  1.  The  cylinder  may  be  filled  to 
the  zero  mark,  then  the  amount  removed  read  on  the  scale.  The 


12  APPARATUS  AND  MANIPULATION 

pipette  is  used  to  measure  one  stated  quantity.  When  the 
drawn  out  tip  is  placed  under  the  surface  of  the  liquid  and 
suction  is  applied  to  the  other  end  until  the  bottom  of  the  men- 
iscus of  the  liquid  is  brought  to  the  level  of  the  mark  on  the 
neck,  the  pipette  will  deliver  the  stated  quantity. 

The  burette  is  filled  to  the  zero  mark,  taking  care  that  no 
bubbles  of  air  remain  in  the  parts  about  the  nozzle,  the  desired 
quantity  of  liquid  withdrawn  from  the  nozzle  at  the  bottom, 
and  the  volume  withdrawn  read  on  the  scale.  All  readings 
should  be  taken  from  the  bottom  of  the  meniscus. 

Fill  a  test-tube  with  20  c.c.  of  water  from  the  pipette,  mark 
the  level  of  the  water  on  the  tube  with  a  paper  label  or  file 
mark.  Determine  the  capacity  of  the  tube  up  to  the  given 
mark  by  means  of  the  cylinder  and  the  burette.  Which  gives 
the  most  accurate  measure  of  volume,  and  why? 

3.  Working  with  Glass  Tubing. 

a.  To  cut  glass  tubing  make  a  file  mark  at  the  desired  spot, 
then  grasping  the  tube  with  both  hands  place  the  two  thumb 
nails  directly  behind  the  mark  and  break  by  pressing  out  and 
pulling  at  the  same  time.  The  sharp  edges  of  the  freshly 
broken  glass  must  be  rounded  off  by  holding  it  in  the  hottest 
part  (Exp.  1)  of  the  Bunsen  flame  till  it  begins  to  glow,  then 
allowing  it  to  cool  slowly.  This  is  called  fire  polishing. 

6.  To  make  a  bend  in  glass  tubing  it  is  desirable  to  heat  the 
tubing  uniformly  through  a  distance  of  5  to  8  cm.  (2  to  3 
inches),  so  as  to  avoid  too  great  thickening  of  the  glass  at  the 
bend.  This  is  accomplished  by  holding  the  glass  parallel  with 
the  flame  of  a  fish  tail  burner,  and  just  above  the  inner  cone, 
i.e.,  in  the  hottest  part  of  the  flame.  The  tubing  should  be 
slowly  rotated  in  the  flame  until  it  has  softened,  then  withdrawn 
from  the  flame,  bent  to  the  desired  angle,  and  allowed  to  cool 
slowly.  Fig.  2. 

c.  To  draw  out  a  fine  capillary,  or  a  jet,  the  glass  should  be 
softened  for  only  1  to  2  cm.  of  its  length.  What  flame  is  best 


GLASS  TUBING  13 

adapted  to  this  purpose?  Rotate  the  glass  tube  in  the  flame 
until  softened  uniformly,  then  remove  it  from  the  flame  and 
pull  it  out  to  the  desired  length. 

d.  After  a  little  practice  in  making  bends  and  jets  make  a 
wash  bottle  as  shown  in  Fig.  3.  A  two-hole  rubber  stopper  or 
a  cork  stopper  may  be  used.  If  a  cork  is  used  it  should  be 
softened  by  means  of  a  press  or  cork  roller  and  the  holes  must 


i  ii 

FIG.  2.  FIG.  3. 

I.  Shows  bends  made  by  heating  the  glass  over  a  very  short  distance,  as 
would  be  done  by  a  Bunsen  burner.  The  glass  buckles  and  is  uneven  in 
thickness.  II.  Shows  similar  bends  made  by  heating  the  glass  over  a 
greater  distance,  using  a  fish-tail  burner.  The  glass  is  even  in  thickness. 


be  made  by  cutting  through  with  a  cork  borer  which  is  slightly 
smaller  in  diameter  than  the  glass  tubing.  Give  the  cork  and 
borer  opposite  rotatory  motion,  with  a  steady  pressure  not  great 
enough  to  tear  the  cork,  until  it  is  cut  half  way  through,  then 
remove  and  start  from  the  other  end,  placing  the  borer  so  that 
the  cuts  will  meet  in  the  middle.  Always  remove  the  cork  from 
inside  the  borer.  If  the  hole  in  the  cork  is  too  small  for  the 
glass  tubing  it  may  be  enlarged  by  using  a  rat-tail  file. 


14  APPARATUS  AND  MANIPULATION 

4.  The  Balance. 

a.  Under  the  direction  of  an  instructor  examine  the  chemical 
balance.     Sit  directly  in  front  of  the  center  of  the  balance  and 
release  the  stop  with  a  slow,  steady  movement.   If  the  pointer 
does  not  swing  raise  the  stop  and  release  again.     The  swing  of 
the  pointer  should  be  entirely  within  the  graduated  portion  of 
the  ivory  scale.     When  the  pointer  is  swinging  freely,  find  the 
center  of  the  arc  described  in  its  swing.     This  is  the  true  zero 
and  its  position  must  be  determined  for  each  weighing,  since  the 
point  changes. 

b.  Place  on  the  left  hand  pan  of  the  balance  (see  rules  in 
paragraph  c)  one  of  the  hard  glass  tubes  which  has  been  previ- 
ously cleaned  and  dried  (see  page  4).     With  the  aid  of  the 
forceps  place  on  the  right  hand  pan  that  combination  of  weights 
which  is  nearest  to,  but  less  than,  the  weight  of  the  tube.     Al- 
ways begin  with  the  largest  weight  and  work  down.     When  the 
tube  is  nearly  counterpoised  close  the  balance  case  and  adjust 
the  rider  upon  the  scale  beam  until  the  pointer  oscillates  about 
the  true  zero,  as  determined  in  a.     Find  the  weight  of  the  tube 
first  by  adding  together  the  values  of  the  weights  which  have 
been  removed  from  the  set;  and  second  by  adding  the  weights 
themselves.     If  these  two  weights  are  identical,  record  in  the 
note-book.     Notice  the  decimal  system  in  the  weights.     For 
example  if  a  dish  is  balanced  when  the  opposite  pan  contains 
10  grams,  2  grams,  500  mg.,  200  mg.,  100  mg.,  10  mg.,  and  the 
rider  is  upon  2  on  the  beam,  the  weight  of  the  dish  is  12.812 
grams.     Now  remove  the  weights,  replacing  each  in  its  proper 
position  in  the  set,  raise  the  rider  from  the  beam,  and  close 
the  balance  case. 

c.  Always  observe  the  following  rules  in  weighing: 

(1)  Handle  the  balance  with  care. 

(2)  Raise  the  stop  gently,  and  always  as  the  pointer  is  passing 
the  zero  mark. 

(3)  Never  add  anything  to  the  pans,  or  remove  anything 
from  them,  while  they  are  swinging. 

(4)  Never  place  chemicals  directly  upon  the  pan.     Use  a 


THE  BALANCE  15 

piece  of  glazed  paper  or  a  watch  glass  which  has  been  properly 
counterpoised. 

(5)  Always  use  forceps  in  handling  the  weights. 

(6)  Place  the  object  to  be  weighed  and  the  weights  as  near 
the  center  of  the  pan  as  possible. 

(7)  Do  not  allow  the  pans  to  swing. 

(8)  In  case  any  solid  is  spilled  in  the  balance  case  use  the 
camels  hair  brush  to  remove  it.     Keep  the  balance  and  the  shelf 
clean. 

(9)  Keep  both  eyes  open  while  reading  the  swing  of  the 
pointer. 

(10)  If  the   balance  is  not  in   proper  adjustment,  call  an 
instructor. 


CHAPTER  II 

PURE  SUBSTANCES  AND  MIXTURES 

5.  Boiling  Point  as  a  Test  for  Purity. 

This  test  depends  on  the  fact  that  a  pure  substance  has  a 
fixed  boiling-point,  while  a  mixture  of  two  or  more  substances 
has  a  boiling-point  which  changes  as  the  boiling  proceeds.  In 
such  a  case  the  substance  which  has  the  lower  boiling-point 
boils  off  in  greater  proportion  at  first,  and  if  the  vapors  are  con- 


FIG.  4. 

densed  a  liquid  can  be  obtained  which  contains  a  higher  per- 
centage of  the  low  boiling  substance  than  the  original  mixture. 
This  fact  is  utilized  in  many  commercial  processes,  as  in  the 
manufacture  of  alcohol. 

a.  Connect  a  50  c.c.  distilling  flask  with  a  Liebig  condenser. 
Fig.  4.     Force  a  slow  stream  of  cold  water  through  the  jacket 

16 


BOILING  POINT  17 

of  the  condenser  in  the  direction  of  the  arrows.  Why  is  the 
water  circulated  in  this  direction?  Half  fill  the  flask  with  dis- 
tilled water  and  stopper  with  a  one-hole  cork  carrying  a 
thermometer.  The  bulb  of  the  thermometer  should  be  just  a 
little  below  the  level  of  the  delivery  tube  of  the  flask.  Why? 
Heat  the  distilling  flask  with  a  small  flame  so  that  the  water 
distils  slowly.  Collect  the  distillate  in  a  small  flask.  After  a 
few  cubic  centimeters  of  water  have  been  collected  read  the 
temperature.  Does  this  stay  constant  as  the  distillation  pro- 
gresses? Is  it  the  temperature  given  in  the  text-book  as  the 
boiling-point  of  water?  If  not,  give  the  reason  for  the  differ- 
ence. What  ought  to  be  specified  when  a  boiling-point  is  given? 

b.  Add  5  c.c.  of  alcohol  to  20  c.c.  of  water  and  repeat  Exp.  a 
with  this  mixture.     What  is  its  boiling-point?     Do  alcohol  and 
water  form  a  mixture  or  a  pure  substance?     Collect  the  first 
few  cubic  centimeters  of  the  distillate  in  an  evaporating  dish  and 
apply  a  flame  to  the  liquid.     In  a  like  manner  collect  a  few 
cubic  centimeters  after  about  half  of  the  original  mixture  has 
been  distilled,  and  als.o  the  last  few  cubic  centimeters.     Do  not 
heat  the  distilling  flask  to  dryness.     Test  each  portion  of  the 
distillate  writh  the  flame.     Account  for  the  difference.     Is  there 
any  alcohol  in  the  last  portion  of  the  distillate?     Is  there  any 
water  in  the  first  portion? 

c.  Dissolve  a  few  small  crystals  of  copper  sulfate  (CuS04)  in 
25  c.c.  of  water  and  distil  the  mixture  until  about  5  c.c.  have  been 
collected.     What   is  the  distillate?    From  what  kind  of  im- 
purities can  water  be  freed  by  distillation?     Evaporate  a  little 
of  the  distillate  on  a  watch  glass,  holding  the  watch  glass  about 
6  inches  above  the  tip  of  the  bunsen  flame.     Evaporate  tap 
water  similarly.     Why  is  distilled  water  used  in  the  laboratory? 

6.  Melting-point  as  a  Test  for  Purity. 

This  test  depends  on  the  fact  that  a  pure  substance  has  a 
fixed  melting-point,  while  a  mixture  of  two  substances  begins 
to  melt  at  a  temperature  different  from  that  at  which  it  is 
completely  melted,  or,  in  reverse  order,  begins  to  solidify  at  a 


18  PURE  SUBSTANCES  AND  MIXTURES 

temperature  different  than  that  at  which  it  is  completely 
solidified.  A  familiar  example  is  found  in  the  melting  of  ice. 
Pure  ice  melts  at  0°  C.,  and  ice  and  water  can  exist  together 
only  at  0°  C.  If  salt  (an  impurity)  is  added  to  water,  the  water 
begins  to  freeze  at  some  temperature  below  0°  C.,  and  the  mass 
is  not  completely  solid  (frozen)  until  —  23°  C.  is  reached.  To 
illustrate  this  test  proceed  as  follows: 

a.  Place  about  2  grams  of  powdered,  hydrated  calcium  chloride 
(CaCl2'6H20)  in  a  clean,  dry  test-tube  and  heat  very  gently 
(this  salt  has  a  very  low  melting-point)  until  the  substance 
softens  enough  so  that  the  bulb  of  a  thermometer  may  be 
inserted  into  it.     Continue  warming,  while  stirring  with  the 
thermometer,  until  a  clear  liquid  is  obtained.     Now  let  the 
melted  mass  cool  gradually,  while  continuing  the  stirring,  by 
holding  the  test-tube  under  running  water,   and  watch  the 
thermometer  constantly.     When  the  solid  particles  first  begin 
to  appear  note  the  temperature,  and  continue  the  cooling  till 
the  contents  of  the  tube  is  solid.     Has  the  temperature  changed 
during  the  process  of  solidification?     Has  the  substance  a  fixed 
freezing-point?     Is  it  a  pure  substance?     If  the  first  attempt 
is  not  successful  this  procedure  may  be  repeated  as  many 
times  as  necessary. 

b.  While  the  thermometer  is  still  in  the  tube  warm  again,  this 
time  noting  the  temperature  at  which  the  substance  melts. 
(Caution:  If  heat  is  applied  too  rapidly  the  melting-point 
cannot  be  observed;  the  heating  should  be  so  gradual  that  the 
solid  is  completely  melted  only  after  several  minutes.)     Is 
the  melting-point  the  same  as  the  freezing-point? 

c.  Now  add  to  the  tube  and  contents  a  small  amount  of  solid 
sodium  chloride  (NaCl)  and  proceed  as  above  to  determine  the 
freezing-point.     Has  the  substance  a  fixed  freezing-point  or 
melting-point  now?     Account  for  the  difference. 

7.  Elements  and  Compounds. 

a.  Grind  4  grams  of  sulfur  and  7  grams  of  iron  filings  to- 
gether in  a  mortar.     Examine  the  mass  with  a  lens.     Place  a 


ELEMENTS  AND  COMPOUNDS  19 

small  amount  of  the  mixture  in  each  of  three  test-tubes.  To  the 
first  add  15  c.c.  of  water  and  shake  thoroughly.  To  the  second 
add  10  c.c.  of  dilute  hydrochloric  acid.  To  the  third  add  5  c.c. 
of  carbon  disulfide,  shake,  filter1  (use  a  dry  filter)  the  liquid 
into  a  watch  crystal  and  allow  it  to  evaporate  spontaneously. 
(Carbon  disulfide  is  inflammable.  Do  not  bring  it  near  a 
flame.)  Describe  what  takes  place  in  each  case. 

b.  Heat  the  remainder  of  the  original  mixture  in  a  test-tube 
until  it  begins  to  glow  strongly.     Allow  the  mass  to  cool,  reduce 
it  to  a  powder,  and  repeat  the  tests  made  in  a.     Account  for 
the  differences  in  the  results  obtained. 

c.  Heat  a  small  portion  of  mercuric  oxide  in  a  dry,  hard  glass 
test-tube.     Watch  for  the  appearance  of  the  evolution  of  a  gas, 
and  test  it  with  a  glowing  splinter.     Note  the  change  in  the 
appearance  of  the  residue  in  the  test-tube,  and  the  deposit  on 
the  cool  parts  of  the  glass.     Is  mercuric  oxide  an  element  or 
compound?     Why?     Have  you  any  experimental  proof  for 
saying  whether  the  products  of  this  decomposition  are  elements 
or  compounds? 

d.  Heat  a  small  portion  of  copper  oxalate  in  the  same  way. 
Test  the  gas  evolved  with  a  glowing  splinter,  and  note  the  change 
in  appearance  of  a  rod  wet  with  lime  water  and  held  at  the 
mouth  of  the  tube.     Is  the  same  gas  evolved  as  before?     What 
remains  in  the  tube?     Continue  the  heating  until  the  residue 
in  the  tube  shows  a  further  change.     Is  this  change  accompanied 
by  the  evolution  of  a  gas? 

e.  Define  an  element,  and  a  compound.   Show  which  materials 
used  above  have  been  proved  to  be  compounds.     If  the  weight 
of  the  residue  from  the  copper  oxalate  had  been  determined 
before  and  after  the  final  change  from  the  dark  red  to  the  black 
substance,  the  change  could  have  been  shown  to  be  accom- 

1  To  prepare  a  filter  paper  fold  it  through  the  middle,  then  fold  the  straight 
edge  onto  itself,  pinching  the  cone  so  made  together  just  at  the  fold.  Now 
open  the  cone  so  that  there  are  three  thicknesses  of  paper  on  one  side  and  one 
on  the  other;  fit  into  the  funnel,  taking  care  that  the  paper  lies  against  the 
glass  wall  on  all  sides,  finish  creasing  the  second  fold  in  the  paper,  and 
moisten  so  that  it  will  stick  in  place. 


20  PURE  SUBSTANCES  AND  MIXTURES 

panied  by  increase  in  weight.     Is  it  then  a  decomposition  or  a 
combination? 

8.  Law  of  Definite  Proportion.     Quantitative. 

Read  the  Directions  for  Quantitative  Experiments  (page  3) 
before  proceeding  with  this  experiment. 

a.  Weigh  accurately  the  two  hard  glass  ignition  tubes  or 
use  those  weighed  in  Exp.  4.  In  the  smaller  place  about  0.5 
gram,  in  the  larger  about  1  gram  of  pure,  dry  potassium 
chlorate  (KC103)  (balance  room)  and  weigh  again.  The  differ- 
ence between  the  two  weights  is  the  exact  weight  of  the 
potassium  chlorate.  Heat  each  tube  gently  in  the  Bunsen 
flame.  The  potassium  chlorate  will  melt,  and  then  appear  to 
boil.  During  the  boiling  the  heating  should  be  just  sufficient 
to  sustain  boiling,  as  too  great  heat  will  cause  the  loss  of  solid 
particles  along  with  the  escaping  gas.  A  glowing  splinter  should 
be  applied  to  the  gas  escaping  from  the  tube,  the  result  to  be 
explained  later.  After  the  boiling  has  ceased  the  tubes  should 
be  heated  with  the  hottest  part  of  the  Bunsen  flame  till  the 
material  is  completely  melted,  then  allow  them  to  cool  and  w^eigh 
again.  Repeat  the  heating  and  weighing  until  constant  weight 
is  obtained.  The  difference  between  the  last  weight  and  that 
of  the  empty  tube  gives  the  weight  of  the  solid  residue  resulting 
from  the  decomposition  of  potassium  chlorate  by  heat.  Keep 
these  tubes  with  tfie  residues  until  your  results  have  been  approved 
by  an  instructor. 

6.  Calculate  the  per  cent  of  residue  in  each  case  and  compare. 
State  the  law  of  definite  proportion  and  show  how  this  experi- 
ment supports  the  law. 

c.  This  example  is  a  case  of  the  decomposition  of  a  more  com- 
plex compound  (potassium  chlorate)  into  an  element  (oxygen) 
and  a  simpler  compound  (potassium  chloride).  Write  the  equa- 
tion. How  was  the  oxygen  recognized?  Dissolve  the  residue, 
after  the  results  of  the  quantitative  work  have  been  approved,  and 
test  the  solution  with  a  solution  of  silver  nitrate  (AgN03). 
Apply  the  same  test  to  a  solution  of  some  of  the  original  po- 
tassium chlorate.  Interpret  the  results. 


CHAPTER  III 

OXYGEN 

9.  Catalysis. 

a.  Melt  a  small  sample  of  potassium  chlorate  (KC103)  in  a 
dry  test-tube  and  note  the  rate  at  which  decomposition  proceeds; 
then  drop  into  the  melted  mass  a  pinch  of  manganese  dioxide 
(Mn02)  (known  to  be  free  from  organic  matter)  and  note  the 
change  in  the  rate  of  decomposition.  Allow  the  mass  to  cool 
and  dissolve.  Has  there  been  any  apparent  change  in  the 
manganese  dioxide? 

6.  Place  as  much  sodium  peroxide  (Na202)  as  can  be  piled  on 
a  dime  in  each  of  two  porcelain  crucibles;  to  one  of  the  crucibles 
add  a  very  small  amount  of  copper  oxide  (CuO),  then  fuse  the 
material  in  each  crucible,  allow  to  cool  and  to  each  add  water  and 
note  the  rate  of  evolution  of  gas.  Test  the  gas  evolved  with  a 
glowing  splinter.  What  is  it?  Does  the  copper  oxide  have  any 
effect  on  the  nature  of  the  chemical  change?  On  the  speed? 
Has  it  undergone  any  change  during  the  reaction?  Boil  the 
mixture  in  the  crucible  containing  the  copper,  and  filter.  The 
copper  oxide  (CuO)  was  converted  into  the  hydroxide  (Cu(OH)2 
or  CuO-H20)  by  the  fusion,  and  on  boiling  changed  back  to 
copper  oxide. 

10.  Preparation  and  Properties  of  Oxygen. 

a.  Calculate  from  the  data  of  Exp.  8,  and  from  the  ca- 
pacity of  a  gas-collecting  bottle  and  the  weight  of  1  liter 
of  oxygen,  the  amount  of  potassium  chlorate  (KC103)  re- 
quired to  give  oxygen  enough  to  fill  four  of  the  gas-collecting 
bottles.  Weigh  out  the  calculated  amount  of  potassium 
chlorate  on  the  rough  laboratory  balance,  mix  with  one-fourth 
its  volume  of  manganese  dioxide  (Mn02)  and  place  in  the  large 

21 


22  OXYGEN 

hard  glass  test-tube  which  should  be  fitted  with  a  stopper 
carrying  an  outlet  tube,  and  which  must  be  proved  to  be  air-tight 
before  proceeding.  A  small  portion  of  this  mixture  should  be 
tested  by  heating  in  another  tube.  If  the  reaction  takes  place 
quietly  it  is  safe  to  conclude  that  organic  matter  is  absent  and 
the  reaction  will  proceed  quietly  in  the  large  tube.  Connect 
the  delivery  tube  with  a  glass  tube  leading  to  the  trough  in 
which  the  gas-collecting  bottles  have  been  filled  with  water 
and  inverted  (Fig.  5)  and  gently  heat  the  potassium  chlorate 
till  gas  is  evolved,  collecting  the  gas  in  the  inverted  bottles. 
After  the  bottles  have  been  filled  with  gas  a  glass  plate  should 


FIG.    5. 

be  placed  over  the  mouth  while  still  under  water,  then  the 
bottle  and  cover  may  be  removed  and  set  aside  for  future  use. 
Always  withdraw  the  delivery  tube  from  below  the  surface  of 
the  water  before  allowing  the  generating  tube  to  cool.  Why? 
The  supply  of  the  gas  may  be  stopped  almost  instantaneously 
at  any  time  by  withdrawing  the  heat  from  the  generating  tube. 
Test  the  four  bottles  of  the  gas  as  follows,  covering  the  bottles 
again  as  soon  as  the  spoon  is  withdrawn. 

b.  Into  one  insert  an  iron  wire  or  picture  cord,  the  tip  of 
which  has  been  heated  and  dipped  into  sulphur,  then  lit. 

c.  Into  the  next  insert  some  burning  sulphur  on  the  iron 
spoon. 


DENSITY  OF  OXYGEN  23 

d.  Into   the  third   insert  a  spoon   carrying  some   burning 
red  phosphorus. 

e.  Into  the  fourth  insert  some  burning  magnesium  ribbon, 
held  with  the  pincers. 

/.  Add  about  10  c.c.  of  water  to  each  of  the  last  three  bottles 
and  shake,  then  test  the  solutions  with  both  blue  and  red 
litmus  paper,  testing  at  the  same  time  a  sample  of  the  water 
added.  What  has  been  the  nature  of  the  reaction  in  these 
cases?  Why  did  combustion  cease  before  all  the  material  had 
burned?  Are  the  products  of  combustion  soluble  in  water? 
To  what  can  you  ascribe  the  difference  in  the  effect  on  litmus 
paper?  Classify  these  three  elements  according  to  the  effect 
of  the  solution  of  their  oxides  on  litmus.  Are  they  at  the  same 
time  classified  with  respect  to  any  physical  properties?  Why  was 
it  necessary  to  test  a  small  portion  of  the  mixture  of  potassium 
chlorate  and  manganese  dioxide  by  heating  separately?  What 
would  be  the  result  if  combustible  matter,  as  charcoal,  had  been 
present  in  the  mixture? 

11.  The  Effect  of  Concentration  on  the  Speed  of  a  Reaction. 

Will  iron  wire  burn  in  the  air  as  it  does  in  oxygen?  Burn 
sulfur,  phosphorus  and  magnesium  in  bottles  filled  with  air. 
Note  the  rate  of  combustion  as  compared  with  burning  in 
oxygen  and  account  for  the  difference.  Add  water  and  test 
with  litmus  as  in  Exp.  10  /.  Are  the  products  of  the  reaction 
the  same  in  the  two  cases? 

12.  To  Determine  the  Weight  of  a  Liter  of  Oxygen.     Quan- 
titative. 

Various  methods  for  determining  the  density,  or  weight  per 
liter  of  a  gas,  may  be  used.  In  the  following  experiment  the 
weight  of  the  gas  will  be  determined  by  the  loss  in  weight  of  a 
material,  due  to  evolution  of  the  gas  in  question,  while  the 
volume  of  the  gas  will  be  measured  directly.  Having  the 
weight  of  a  given  volume  the  weight  per  liter  can  then  be  calcu- 
lated by  direct  proportion. 


24 


OXYGEN 


a.  Fit  the  large  ignition  tube  with  a  rubber  stopper  and  de- 
livery tube  bent  so  that  it  can  be  connected  with  a  gas  burette 
(Fig.  6).  Fill  the  burette  with  water  and  see  that  there  is  no  air 
in  the  connecting  rubber  tubing  by  raising  and  lowering  the 
reservoir  (R)  several  times,  while  the  pinch  cock  is  open.  Con- 
nect the  burette  with  the  ignition  tube  and  test  for  air  tightness. 
Then  remove  the  ignition  tube,  weigh  it,  add  0.2  to  0.5  gram  of 
potassium  perchlorate  (KC104)  and  weigh  again.  Bring  the 
level  of  the  water  in  the  burette  near  the  zero  mark,  attach  the 


FIG.   6. 


ignition  tube  and  open  the  cock.  If  the  level  of  the  water 
in  the  burette  remains  constant  for  two  minutes  (showing  the 
apparatus  to  be  air  tight)  bring  the  gas  to  atmospheric  pressure 
by  raising  or  lowering  the  reservoir  until  the  level  of  the  water 
in  the  two  tubes  is  the  same,  then  record  the  burette  reading. 
Now  heat  the  potassium  perchlorate  gently  until  about  30  c.c. 
of  gas  has  been  evolved,  then  allow  to  stand  till  the  tube  has 
reached  room  temperature.  Read  the  volume  of  the  gas  in 
the  burette  after  bringing  the  level  of  the  water  in  the  reservoir 
even  with  that  in  the  burette  (the  cock  should  remain  open 


DENSITY  OF  OXYGEN  25 

through  all  this  procedure),  read  the  temperature  and  baro- 
metric pressure  in  the  laboratory,  then  remove  and  weigh  the 
ignition  tube  and  residue.  If  moisture  has  collected  in  the 
ignition  tube  allow  it  to  stand  over  night,  or  until  it  has  com- 
pletely dried  by  spontaneous  evaporation,  before  weighing. 
The  difference  between  the  two  burette  readings  is  the  volume 
of  oxygen,  the  difference  between  the  weights  of  the  tube  and 
potassium  perchlorate  and  the  tube  and  residue  is  the  weight  of 
the  oxygen. 

b.  Since  the  volume  of  a  gas  varies  greatly  under  varying  con- 
ditions of  temperature  and  pressure,  it  is  convenient  to  state  the 
weight  of  a  given  volume  under  " standard"  conditions,  i.e., 
at  0°  C.  and  760  mm.  pressure.     Calculate  the  volume  the 
oxygen  would  occupy  under  standard  conditions,  remembering 
that  the  gas  measured  in  the  burette  is  a  mixture  of  oxygen 
and  water  vapor,  and  that  the  atmospheric  pressure  read  on 
the  barometer  must  be  corrected  for  temperature  (see  page  5). 
The  pressure  correction  to  be  applied  for  both  water  vapor  and 
the  temperature  of  the  barometer  may  be  found  in  tabulated 
form  on  page  6.     From  the  weight  of  the  given  volume  of 
oxygen  calculate  the  weight  of  one  liter.     Calculate  the  per 
cent  by  which  your  result  differs  from  the  accepted  value, 
1.429  grams.     Keep  the  ignition  tube  and  residue  for  further 
work.     (Exp.  24.) 

c.  Why  is  it  necessary  to  allow  for  the  vapor  pressure  in  the 
gas?     Why  does  not  the  fact  that  the  ignition  tube  was  full 
of  air,  some  of  which  has  gone  with  the  oxygen  into  the  gas 
burette,  vitiate  the  results?     How  would  it  affect  the  result 
if  the  ignition  tube  had  not  cooled  completely  before  reading 
the  volume  of  oxygen?     Why  must  the  water  in  the  reservoir 
be  brought  to  the  level  of  that  in  the  burette  before  reading 
the  volume  of  oxygen? 


CHAPTER  IV 

HYDROGEN 
13.  Preparation  and  Properties  of  Hydrogen. 

a.  Construct  an  apparatus  consisting  of  a  250  c.c.  generating 
flask  (the  Erlenmeyer  is  well  adapted  to  this  purpose)  stoppered 
with  a  cork  carrying  a  thistle  tube  reaching  to  within  0.5  cm. 
of  the  bottom  of  the  flask  and  a  delivery  tube  bent  for  the 


FIG.    7. 

delivery  of  gas  into  bottles  inverted  in  the  pneumatic  trough 
(Fig.  7).  This  apparatus  must  be  air  tight  and  may  best  be 
tested  for  tightness  by  closing  the  open  end  of  the  delivery 
tube  with  a  rubber  tube  and  pinch  cock,  then  pouring  water 
into  the  thistle  tube.  Explain  why  the  water  cannot  run  down 
into  the  flask  if  the  apparatus  is  tight. 

If  the  following  directions  are  observed  carefully  there  will 

26 


PREPARATION  OF  HYDROGEN 


27 


be  no  danger  of  an  explosion,  though  as  further  protection  a 
towel  may  be  wrapped  about  the  generating  flask,  if  desired. 
6.  Place  about  5  grams  of  granulated  zinc  in  the  generating 
flask,  pour  in  enough  dilute  hydrochloric  acid  to  cover  the 
lower  end  of  the  thistle  tube  and  collect,  the  gas  by  holding  an 
inverted  test-tube  over  the  end  of  the  delivery  tube.  Carry  the 
test-tube  mouth  downward  to  the  Bunsen  burner,  which  must 
be  at  least  two  feet  from  the  delivery  tube,  repeating  the  process 
until  the  gas  in  the  test-tube  burns  quietly.  Try  the  effect  if 


FIG.    8. 


the  test-tube  of  gas  is  carried  mouth  upward  to  the  flame. 
What  information  does  this  give  regarding  the  density  of  the  gas? 
Collect  a  bottle  of  hydrogen  by  displacement  of  water  and 
while  holding  it  mouth  downward  insert  into  it  a  burning 
splinter.  Why  does  the  flame  go  out,  and  why  is  the  splinter 
ignited  again  as  it  is  withdrawn  from  the  tube  ? 

c.  Attach  a  glass  jet  to  the  delivery  tube  by  means  of  a  short 
piece  of  rubber  tubing.  Collect  a  test-tube  of  hydrogen  from 
this  jet  by  displacement  of  air  (should  the  test-tube  be  inverted 
or  upright?)  and  after  lighting  the  test-tube  of  gas  at  the  Bunsen 


28  HYDROGEN 

flame  carry  it  back  quickly  to  the  jet.  If  the  hydrogen  is 
pure  enough  to  be  lighted  with  safety  it  will  light  in  this  way. 
Never  apply  a  match  to  a  hydrogen  generator,  but  always  light  it 
with  a  test-tube  of  the  burning  gas  as  described.  Hold  a  watch 
glass  above  the  flame  for  a  moment.  What  is  the  deposit? 
Explain  its  formation. 

d.  Remove  the  glass  jet  and  replace  it  with  a  hard  glass  tube 
open  at  each  end  and  containing  about  0.5  gram  copper  oxide 
(CuO),  Fig.  8.     After  again  testing  the  apparatus  to  make  sure 
it  is  free  from  an  explosive  mixture  warm  the  copper  oxide 
gently  while  passing  hydrogen  over  it.     Hold  a  watch  glass 
at  the  outer  end  of  the  combustion  tube.     What  is  deposited? 
What   remains   in   the    combustion   tube?     What   has    been 
reduced  during  this  action?     What  has  been  oxidized?     Write 
the  reaction. 

e.  Clean  out  the  hard  glass  tube  and  repeat  d  using  magnetic 
oxide  of  iron  in  place  of  cupric  oxide,  and  observing  all  pre- 
cautions to  avoid  an  explosion.     Write  the  equation  for  this 
reaction.     What  has  been  oxidized  during  this  action? 

/.  May  other  metals  and  acids  be  used  in  the  preparation  of 
hydrogen?  Test  this  by  pouring  dilute  sulfuric  acid  on  sam- 
ples of  magnesium,  aluminium  and  copper,  each  in  a  separate 
test-tube.  The  gas  evolved  may  be  tested  by  bringing  the 
mouth  of  the  test-tube  to  the  flame.  Add  a  few  drops  of 
dilute  nitric  acid  to  a  sample  of  lead  in  a  test-tube.  Does  the 
gas  evolved  have  the  properties  of  hydrogen?  In  all  these 
cases  if  the  reaction  is  not  spontaneous  the  test-tube  and  its 
contents  should  be  warmed  gently  before  drawing  any 
conclusions. 

14.  The  Rate  of  Diffusion  of  Gases. 

Draw  out  the  closed  end  of  a  test-tube  to  a  capillary  (if  the 
closed  end  gets  too  hot  to  hold,  a  piece  of  glass  tubing  may  be 
sealed  to  it  by  heating  the  two  together  for  a  moment)  cut  the 
capillary,  and  reduce  the  size  of  the  opening  by  heating,  gently, 
until  when  the  tube  full  of  air  is  inserted  open  end  down  into  a 


DIFFUSION  OF  GASES 


29 


beaker  full  of  water  (Fig.  9)  it  takes  about  one-half  minute  for  the 
water  to  rise  in  the  tube  to  a  mark  made  about  two  inches 
from  the  open  end.  The  water  in  the  beaker  must  of  course 
come  above  this  mark  on  the  test-tube.  In  order  that  the  fol- 
lowing measurements  may  be  comparable  the  test-tube  should 
always  be  brought  to  the  bottom  of  the  beaker,  tipped  just 
enough  so  there  is  free  access  of  water  from  the  outside,  and 
the  level  of  the  water  in  the  beaker  should  be  always  the 
same.  Now  determine  accurately  the  time  elapsing  between 
the  moment  when  the  tube  is  inserted  into  the  beaker  of  water 
and  when  the  water  comes  to  the  mark  on  the  test- tube, 
repeat  this  several  times  with  the  tube  full  of  air,  recording 
each  determination,  then  fill  the  tube  with 
hydrogen  by  displacement  of  air  and  again 
determine  several  times  the  time  required 
for  the  water  to  rise  to  the  mark,  as  before. 
Remember  that  in  order  to  completely  fill 
this  tube  with  the  easily  diffusible  hydro- 
gen a  rapid  stream  of  the  gas  should  be 
led  into  it,  and  in  order  to  keep  the  hydro- 
gen after  the  tube  is  filled  the  thumb  may 
be  placed  over  the  open  end,  but  the  time 
elapsing  between  filling  the  tube  and  insert- 
ing it  into  the  water  must  be  as  brief  as  pos- 
sible. Repeat,  filling  the  tube  with  carbon  dioxide  from  a 
Kipp  generator,  observing  all  the  precautions  as  in  the  case  of 
hydrogen.  Calculate  the  densities  of  these  gases  from  their 
rates  of  diffusion,  taking  hydrogen  as  1,  and  compare  with  the 
values  given  in  the  table. 

15.  The  Combining  Weight  of  Copper.     Quantitative. 

a.  Fit  up  the  hydrogen  generator  as  described  in  Exp.  13,  and 
connect  the  delivery  tube  with  a  hard  glass  combustion  tube 
open  at  each  end,  held  horizontally  by  the  iron  clamp,  and  from 
which  a  right  angle  bend  leads  to  the  bottom  of  a  test-tube, 
Fig.  8.  After  proving  the  apparatus  to  be  air  tight  disconnect 


FIG.  9 


30  HYDROGEN 

the  combustion  tube,  place  in  it  about  0.5  gram  (do  not  weigh 
accurately)  of  powdered  copper  oxide  (CuO),  and  again  connect 
with  the  generator.  Place  about  7  grams  of  granulated  zinc  in 
the  generator  and  add  dilute  hydrochloric  acid  till  the  lower 
end  of  the  thistle  tube  is  covered.  See  that  the  delivery  tube  is 
at  least  two  feet  away  from  any  flame.  Collect  a  test-tube  of  the 
gas  by  displacement  of  air,  and  carry  it,  inverted,  to  the  flame, 
repeating  the  process  until  the  gas  in  the  test-tube  burns 
quietly.  Why?  Now  heat  the  copper  oxide  gently,  having  the 
delivery  tube  bent  down  and  leading  to  the  bottom  of  a  test- 
tube.  Notice  the  change  in  the  copper  oxide,  and  continue 
the  heating  until  no  further  change  takes  place  (this  will  require 
about  20  minutes),  and  all  the  water  formed  has  been  driven 
out  of  the  combustion  tube  into  the  test-tube,  then  allow  the 
tube  and  contents  to  cool  while  hydrogen  is  still  passing  through. 
(Be  sure  that  the  evolution  of  hydrogen  continues  through  all 
the  heating.)  Write  the  equation  for  the  reaction.  What  is 
the  reducing  agent  in  this  action?  What  is  the  oxidizing  agent? 

6.  If  the  copper  oxide  has  been  completely  reduced,  and  the 
product  resulting  allowed  to  cool  in  an  atmosphere  of  hydrogen, 
it  should  be  pure  copper.  A  weighed  amount  of  this  copper 
may  now  be  converted  to  the  oxide,  weighed,  and  from  these 
two  weights  the  weight  of  oxygen  in  the  oxide  is  obtained  by 
difference.  The  combining  weight  of  copper,  or  the  weight  of 
copper  which  will  combine  with  8  grams  of  oxygen,  may  be 
calculated  by  proportion. 

In  order  to  do  this,  clean  and  weigh  a  porcelain  crucible. 
Place  in  it  part  of  the  copper  obtained  above,  and  weigh  again. 
Now  heat  the  crucible  and  contents  until  the  red  copper  color 
has  entirely  disappeared,  stirring  occasionally  with  the  platinum 
wire.  Allow  the  crucible  to  cool,  and  weigh  it.  Heat  it  again 
and  weigh  again,  until  two  successive  weighings  are  within 
2  mg.  the  same.  Calculate  the  combining  weight  of  copper. 

16.  The  Equivalent  Weight  of  Magnesium,  Exp.  93,  page  98, 
may  be  Performed  here  if  Desired. 


CHAPTER  V 

WATER  AND  HYDROGEN  PEROXIDE 
17.  Composition  of  Water. 

a.  The  qualitative  composition  of  water  by  synthesis  has 
been  illustrated.  Review  the  experiments. 

6.  The  qualitative  composition  of  water  by  analysis  may  be 
illustrated  by  the  decomposition  of  water  vapor  by  hot  iron  filings. 

Fit  a  100  c.c.  flask  with  a  stopper  carrying  a  thistle  tube  and 
an  obtuse  angle  delivery  tube,  which  should  be  connected  with  a 


FIG.  10. 

hard  glass  combustion  tube,  the  further  end  of  the  combustion 
tube  being  connected  with  a  delivery  tube  leading  under 
water,  Fig.  10.  Make  all  connecting  tubes  as  short  as  possible 
so  as  to  avoid  condensation  of  the  steam.  After  making  sure 
that  the  apparatus  is  air  tight  place  about  50  c.c.  of  water  in 
the  flask  and  about  a  gram  of  iron  filings  in  the  combustion  tube. 
Clamp  the  tube  so  that  it  slopes  down  toward  the  flask  slightly 
(to  prevent  water  from  condensing  and  running  over  the  hot 

31 


32  WATER 

part  of  the  glass)  and  heat  the  iron  filings  with  one  Bunsen 
burner,  at  the  same  time  bringing  the  water  in  the  flask  just  to 
boiling  with  another  burner.  For  successful  results  the  iron 
must  be  kept  at  a  low  red  heat  while  steam  is  passing  slowly 
over  it.  Collect  the  decomposition  product  of  the  steam  by 
displacement  of  water,  and  test  for  hydrogen.  What  was 
the  source  of  the  hydrogen?  Where  will  you  expect  to  find 
the  oxygen?  Remove  some  of  the  material  from  the  com- 
bustion tube.  Does  it  appear  the  same  as  the  iron  filings  used? 
Refer  to  the  text-book  for  the  composition  of  the  iron  oxide, 
and  write  the  equation. 

Compare  this  equation  with  the  one  written  in  Exp.  13  e. 
What  is  the  relationship  between  these  two  equations?  What 
kind  of  a  reaction  is  this?  What  causes  the  equilibrium  to 
shift  in  opposite  directions  in  these  two  experiments? 

18.  Solubility  of  Salts  in  Water.     Quantitative. 

The  solubility  of  a  salt  is  defined  as  the  weight  of  salt  in 
100  grams  of  a  saturated  solution  at  the  given  tempera- 
ture. The  solubility  determination  is  usually  made  by  weigh- 
ing a  given  amount  of  solution  known  to  be  saturated  at  a 
given  temperature,  removing  the  water  by  spontaneous 
evaporation  and  finally  by  heating,  then  weighing  the 
dry  salt. 

a.  Cover  each  of  two  evaporating  dishes  with  a  watch  glass  or 
glass  plate  and  weigh  the  dish  with  cover. 

b.  Prepare  a  saturated  solution  of   potassium   dichromate 
(K2Cr207)  or  potassium  chloride  (KC1)  by  adding  the  powdered 
salt  to  about  50  c.c.  of  warm  (40°)  water  until  some  remains 
undissolved  after  stirring  intermittently  for  ten  minutes.     The 
success  of  this  experiment  depends  on  the  preparation  of  a  truly 
saturated  solution,  therefore  plenty  of  time  must  be  allowed  for 
all  of  the  salt  that  will  to  dissolve;  the  temperature  must  be  kept 
as  constant  as  possible,  the  solution  must  be  stirred,  and  the 
salt  must  be  powdered  so  that  it  will  dissolve  rapidly.    Pour 
about  10  c.c.  (estimate  the  quantity,  do  not  measure  it)  of  this 


SUPERSATURATED  SOLUTIONS  33 

saturated  solution  into  one  of  the  weighed  evaporating  dishes, 
noting  the  temperature  of  the  solution  at  the  moment  of  pouring. 
Weigh  the  solution  in  the  covered  dish  immediately,  and  only 
to  the  second  decimal  place,  i.e.,  record  the  weight  as  soon  as 
you  come  to  the  place  where  0.01  gram  more  is  too  much.  Then 
place  the  dish  in  a  safe  place  in  the  desk  and  remove  the  cover, 
leaving  it  to  evaporate  spontaneoulsy,  or  hasten  the  evaporation 
by  placing  the  dish  on  the  steam  bath.  In  the  meantime 
the  saturated  solution  will  have  cooled  down,  and,  after 
stirring  vigorously  and  allowing  the  crystals  to  settle,  a 
second  sample  should  be  placed  in  the  second  weighed  dish, 
the  temperature  at  the  time  of  pouring  noted,  and  the  pro- 
ceeding repeated. 

c.  After  the  solutions  in  the  dishes  appear  to  have  evaporated 
to  dryness  replace  the  covers  and  heat  each  one  gently  with  the 
Bunsen  burner,  keeping  the  cover  on  to  avoid  loss  from  decrepi- 
tation of  the  crystals.     Decomposition  will  result  if  potassium' 
dichromate  is  heated  above  its  melting-point.     If  moisture 
collects  on  the  cover  warm  until  it  has  entirely  disappeared, 
then  allow  to  cool  and  weigh. 

d.  From  the  data  calculate  the  weight  of  potassium  dichro- 
mate or  potassium  chloride  which  will  dissolve  in  100  grams  of 
water  at  the  two  temperatures.     Look  up  the  solubility  of  the 
salt  used  at  10°  and  20°,  and  draw  a  solubility  curve  from  these 
four  points. 

19.  Supersaturated  Solutions. 

a.  Define  a  solvent,   a  solute,   saturated,   unsaturated  and 
supersaturated   solutions.     Could  a  solution  in  contact  with 
the  solid  solute  remain  permanently  unsaturated  or  supersatu- 
rated?    How  would  you  make  a  supersaturated  solution  of  a 
salt  more  soluble  in  hot  water  than  in  cold?     Of  a  salt  more 
soluble  in  cold  water  than  in  hot  ?     Give  two  examples  of  unsta- 
ble equilibrium  and  show  in  what  respect  each  differs  from 
stable  equilibrium. 

b.  To  25  c.c.  of  water  add  20  grams  of  hydrated  sodium  sul- 
fate  (Na2S04'10H20)  and  warm  with  stirring  until  solution  is 


34  WATER 

complete.  Decant  into  a  perfectly  clean  beaker,  cover  with  a 
watch  glass  and  set  aside,  allowing  the  solution  to  cool  spon- 
taneously. Agitating  the  solution,  or  the  presence  of  any  solid 
matter,  may  cause  crystallization  to  take  place  during  the  cool- 
ing. When  the  solution  has  cooled  to  room  temperature,  if  no 
crystallization  has  taken  place,  drop  in  one  crystal  of  sodium 
sulfate,  and  explain  the  phenomenon.  If  crystallization 
occurs  during  cooling  the  crystals  will  have  to  be  redissolved 
by  warming  again,  and  the  cooling  and  subsequent  treatment 
with  a  crystal  of  sodium  sulfate  repeated.  What  is  the  condi- 
tion of  the  solution  (saturated,  unsaturated  or  supersaturated) 
(1)  when  poured  into  the  clean  beaker,  (2)  when  the  undisturbed 
solution  has  cooled,  (3)  after  crystallization  has  taken  place? 
What  test  could  be  applied  to  determine  whether  a  given  solu- 
tion is  saturated,  supersaturated  or  unsaturated? 

20.  Water  of  Hydration. 

Heat  1  gram  samples  of  copper  sulfate,  potassium  nitrate, 
sodium  chloride,  sodium  sulfate  and  potassium  dichromate 
each  in  a  separate  test-tube,  and  determine  which  of  them  con- 
tain water  of  hydration.  Place  the  samples  of  copper  sulfate 
and  sodium  sulfate,  after  they  have  been  heated  in  the  test- 
tube,  on  watch  glasses,  and  on  two  other  watch  glasses  place 
samples  of  the  same  materials  which  have  not  been  heated. 
Label  these  four  samples  and  allow  them  to  stand  over  night 
in  the  desk,  then  examine.  In  which  case  has  the  change 
brought  about  by  heating  been  reversed?  To  what  is  this 
reversal  due?  Which  of  the  salts  is  efflorescent  and  which 
deliquescent?  Describe  a  means  by  which  a  deliquescent  salt 
may  be  made  to  effloresce. 

21.  Natural  Waters. 

a.  Evaporate  about  5  c.c.  each  of  distilled  and  tap  water  on 
watch  glasses  on  the  water  bath,  note  the  quantity  of  the  residue 
in  each  case  and  test  it  to  see  if  it  is  again  soluble  in  water. 

b.  Test  portions  of  tap  and  distilled  water  for  chlorides  by 
adding  a  few  drops  of  dilute  C.P.  nitric  acid  (HN03)  (side  shelf) 


HYDROGEN  PEROXIDE  35 

and  silver  nitrate  (AgN03)  solution.  (A  white  precipitate  in- 
dicates the  presence  of  chlorides.)  Test  other  portions  for 
carbonates  by  adding  freshly  filtered  lime  water  (a  white  precipi- 
tate indicates  the  presence  of  carbonates)  and  for  calcium  salts 
by  adding  to  fresh  portions  ammonium  hydroxide  (NH4OH) 
and  ammonium  oxalate  ((NH4)2C204)  solution  (a  white  preci- 
pitate shows  the  presence  of  calcium  salts) .  Compare  the  tests. 
c.  If  rain  water  and  sea  water  are  available  repeat  these  tests 
on  them,  and  compare  all  four  waters. 

HYDROGEN  PEROXIDE 

22.  Preparation  and  Properties  of  Hydrogen  Peroxide. 

a.  To  make  a  solution  of  hydrogen  peroxide  add  gradually 
about  1  gram  of  sodium  peroxide  (Na202)  (carry  this  in  a 
watch  glass  or  evaporating  dish,  as  it  will  cause  spontaneous 
combustion  of  organic  material  such  as  paper)  to  about  50  c.c. 
of  water  in  a  flask,  which  should  be  kept  cool  by  holding  under 
running  water.  Why?  Then  add  dilute  sulfuric  acid,  gradu- 
ally and  with  stirring,  until  litmus  paper  placed  in  the  solution 
just  turns  red,  then  test  the  solution  as  follows  (save  some  of  this 
solution  for  Exp.  23) : 

6.  Prepare  some  lead  sulfide  by  passing  hydrogen  sulfide 
through  1  c.c.  of  a  solution  of  lead  acetate  (Pb(C2H302)2),  or 
by  adding  a  few  drops  of  acetic  acid  to  1  c.c.  of  lead  acetate, 
then  adding  0.5  c.c.  of  ammonium  sulfide.  Add  a  few  cubic 
centimeters  of  the  hydrogen  peroxide  solution  to  this  mixture, 
warm,  and  note  the  change  in  color  of  the  precipitate.  What 
is  the  reaction? 

c.  In  another  test-tube  place  1  c.c.  of  a  solution  of  potassium 
permanganate  (KMnOJ,  add  a  few  drops  of  dilute  sulfuric 
acid,  and  then  the  hydrogen  peroxide.  Test  the  escaping  gas 
with  a  glowing  splinter.  What  is  it?  Write  the  equation. 

23.  To  Determine  the  Concentration  of  the  Hydrogen  Peroxide. 
Quantitative. 

Fill  a  gas  burette  with  water  and  attach  a  short-stemmed 
funnel  to  the  rubber  tubing  above  the  stopcock  (Fig.  11).  Bring 


36 


WATER 


the  water  up  to  the  top  of  the  graduated  tube  and  close  the 
cock,  then  introduce  into  the  funnel  10  c.c.  of  the  hydrogen 
peroxide  solution  prepared  in  Exp.  22,  measured  by  the  pipette. 
By  opening  the  cock  carefully  the  solution  can  be  drawn  into 
the  gas  burette  without  admitting  any  air  after  it.  If  air  is 
admitted  it  must  be  forced  out  by  raising  the  reservoir  and 
opening  the  cock,  but  the  mixing  of  the  solution  with  the  water 
in  the  burette  which  this  causes  is  not  desirable.  Now  place 
about  20  c.c.  of  potassium  permanganate  solution,  which  has 
been  acidified  with  2  c.c.  of  dilute  sulfuric  acid,  in  the  funnel 
and  draw  about  2  c.c.  at  a  time  into  the  burette,  until  the  color 
of  the  potassium  permanganate  in  the  burette  is  permanent, 

remembering  that  the  solution  can  be 
drawn  in  only  when  the  level  of  the 
water  in  the  outer  reservoir  is  below 
that  in  the  burette,  and  do  not  at  any 
time  draw  in  all  the  solution  from  the 
funnel.  When  the  reaction  has  ceased, 
as  indicated  by  the  permanence  of  the 
color,  read  the  volume  of  oxygen  lib- 
erated at  atmospheric  pressure,  the 
barometer  and  the  thermometer.  Cal- 
culate the  weight  of  oxygen  obtained 
(remember  that  it  was  saturated  with 
moisture).  Write  the  equation  for  the 
reaction  involved.  What  per  cent  of 
the  oxygen  liberated  comes  from  the 
hydrogen  peroxide?  From  what  source 
does  the  remaining  oxygen  come? 
Calculate  the  weight  of  the  hydrogen  peroxide  which  was  present 
in  the  10  c.c.  portion  of  the  solution.  Assuming  the  solution 
to  have  a  specific  gravity  of  1,  calculate  the  per  cent  by  weight 
of  hydrogen  peroxide  in  the  solution. 
24.  Law  of  Multiple  Proportion.  Quantitative. 

In  experiment  8  (the  law  of  definite  proportion)  the  amount  of 
potassium   chloride   in   potassium   chlorate   was   determined. 


FIG.  11. 


LAW  OF  MULTIPLE  PROPORTION  37 

The  amount  of  potassium  chloride  in  potassium  perchlorate  may 
be  determined  in  a  similar  fashion,  and  from  these  two  deter- 
minations the  amount  of  oxygen  united  with  a  given  amount  of 
potassium  chloride  in  each  of  the  compounds  calculated.  If 
the  ignition  tube  and  residue  from  Exp.  12  have  been  kept,  the 
residue  (the  partially  decomposed  potassium  perchlorate) 
should  be  heated  until  there  is  no  further  evolution  of  oxygen, 
then  cooled  and  weighed.  From  the  weights  taken  in  Exp.  12. 
and  this  final  weighing,  calculate  the  per  cent  of  potassium 
chloride  in  potassium  perchlorate. 

From  the  data  obtained  in  Exp.  8  calculate  the  amount  of 
oxygen  which  is  combined  with  1  gram  KC1  in  KC103.  From 
the  data  in  this  experiment  calculate  the  amount  of  oxygen 
which  is  combined  with  1  gram  KC1  in  KC104.  What  is  the 
ratio  of  these  two  amounts? 

State  the  law  of  multiple  proportion  and  show  how  this 
ratio  illustrates  it. 


CHAPTER  VI 

THE  HALOGEN  FAMILY 

25.  Preparation  of  Chlorine. 

Caution:  Do  not  allow  chlorine  to  escape  into  the  air  un- 
necessarily, as  the  fumes  are  poisonous  and  cause  very  unpleasant 
colds  if  breathed  in  small  quantities. 

By  what  general  method  may  chlorine  be  prepared  from  hy- 
drochloric acid  ?  What  sort  of  a  reaction  would  this  be  ?  What 
general  class  of  agents  would  be  used  for  this  purpose?  Wliat 
would  you  expect  to  be  the  result  if  materials  which  give 
oxygen  on  heating  were  warmed  with  hydrochloric  acid? 
Test  your  conclusion  by  placing  very  small  amounts  of 
potassium  chlorate  (KC103),  potassium  perchlorate  (KC104), 
and  manganese  dioxide  (Mn02)  in  separate  test-tubes  and 
warming  with  about  1  c.c.  of  dilute  hydrochloric  acid.  The 
gas  evolved  should  be  tested  for  chlorine  by  holding  a  strip 
of  paper  wet  with  a  starch  potassium  iodide  solution  (made 
by  mixing  equal  volumes  of  starch  solution  and  potassium 
iodide  (KI)  solution)  above  the  mouth  of  the  tube.  If  chlor- 
ine is  present  the  paper  will  turn  blue,  due  to  the  formation 
of  a  compound  of  starch  and  iodine  which  is  blue,  but  a  large 
amount  of  chlorine  may  further  change  the  color  to  brown, 
then  bleach  it.  The  brown  color  is  that  of  free  iodine,  and 
the  bleaching  is  due  to  the  oxidation  of  iodine  by  chlorine 
and  water,  forming  iodic  and  hydrochloric  acids.  Write  the 
reactions.  As  soon  as  the  test  has  been  made  pour  the  con- 
tents of  the  tube  down  the  sink,  and  follow  with  plenty  of 
water. 

26.  The  Density  of  Chlorine.     Quantitative.     Poison. 

The  method  used  in  determining  the  weight  of  a  liter  of 
oxygen  cannot  be  used  in  this  case  because  of  the  solubility 

38 


DENSITY  OF  CHLORINE 


39 


of  chlorine  in  water,  and  the  lack  of  a  material  which  can  be 
used  to  give  off  chlorine  as  potassium  perchlorate  gives  oxygen. 
The  method  used  for  chlorine  in  this  experiment  consists  in 
weighing  the  same  volume  of  air  and  of  chlorine  under  known 
conditions,  and  finding  the  volume  weighed  by  replacing  the 
chlorine  with  a  liquid,  whose  volume  can  be  measured.  The 
volume  of  the  chlorine  is  then  equal  to  the  volume  of  the 
liquid  necessary  to  replace  it,  and  its  weight  is  equal  to  the 


FIG.  12. 


weight  of  that  same  volume  of  air  +  the  difference  between 
the  weight  of  the  flask  and  air,  and  the  flask  and  chlorine. 

a.  Carefully  clean  and  dry  a  250  c.c.  receiving  flask,  fit  it  with  a 
cork  stopper  and  weigh  the  flask  and  stopper,  recording  the  atmos- 
pheric pressure  and  temperature  at  the  same  time.  Now  fit 
up  a  chlorine  generator  similar  to  that  used  for  hydrogen, 
taking  special  care  to  have  all  rubber  connections  as  short  as 
possible,  since  chlorine  attacks  the  rubber.  The  delivery  tube 
from  the  generator  should  lead,  through  a  cork,  to  the  bottom 


40  THE  HALOGEN  FAMILY 

of  a  Woulff  bottle,  the  other  neck  of  which  is  fitted  with  a 
cork  carrying  a  delivery  tube  which  projects  just  below  the 
surface  of  the  cork,  and  is  bent  so  that  it  will  reach  to  the 
bottom  of  the  receiving  flask  (Fig.  12).  Place  about  4  grams 
of  manganese  dioxide  (Mn02)  in  the  generating  flask  and  add 
10  c.c.  of  concentrated  hydrochloric  acid,  diluted  with  10  c.c. 
of  water.  The  chlorine  coming  from  this  generator  should  be 
dried  by  bubbling  through  concentrated  sulphuric  acid  in  the 
Woulff  bottle  (a  depth  of  1  cm.  is  sufficient)  and  should  then 
be  led  to  the  bottom  of  the  receiving  flask  to  prevent  diffusion 
into  the  air  as  far  as  possible.  Generate  the  chlorine  by 
warming  gently  and  as  soon  as  a  green  gas  appears  in  the 
generating  flask  introduce  the  delivery  tube  into  the  weighed 
flask,  leaving  the  flask  unstoppered  while  chlorine  passes  in. 
As  soon  as  the  flask  seems  to  be  full  of  chlorine,  as  indicated 
by  the  color  of  the  gas,  remove  the  delivery  tube  and  stopper 
the  flask.  Collect  the  rest  of  the  chlorine  in  three  gas-collecting 
bottles  the  mouths  of  which  are  closed  as  nearly  as  possible 
with  glass  plates,  cover,  and  set  aside  for  Exp.  27. 

b.  Weigh  the  stoppered  flask  containing  the  chlorine,  again 
noting  the  atmospheric  pressure  and  temperature,  remove  the 
cork  and  invert  the  flask  over  a  beaker  containing  225  c.c. 
of  water  and  25  c.c.  of  potassium  hydroxide  (KOH)  solution. 
Do  not  allow  the  flask  filled  with  chlorine  to  stand  any  length 
of  time  before  this  absorption,  as  the  chlorine  is  absorbed  by 
the  cork,  ruining  the  experiment.     The  chlorine  will  dissolve 
in  the  alkaline  solution,  causing  the  liquid  to  be  sucked  up  into 
the  flask.     When  the  reaction  seems  to  be  complete  place  the 
beaker  and  flask,  with  its  mouth  still  under  the  surface  of  the 
solution  in  the  beaker,  in  a  sink  full  of  water,  lower  the  beaker, 
bring  the  level  of  the  liquid  inside  the  flask  even  with  the  water 
outside,  and  while  holding  the  flask  in  this  position  put  in  the 
stopper.     Now  remove  the  flask,  and  measure  the  volume  of 
the  solution  it  contains,  which  is  equal  to  the  volume  of  air 
replaced  by  chlorine. 

c.  The  weight  of  the  air  replaced  by  chlorine  may  be  calculated 


CHLORINE  41 

from  its  volume,  the  temperature  and  pressure  at  the  first 
reading  (assume  that  the  air  was  dry)  and  the  weight  of  a 
liter  of  air  under  standard  conditions.  The  difference  between 
the  weight  of  that  volume  of  air  and  of  chlorine  may  be  found 
by  subtracting  the  weight  of  the  flask  and  air  from  the  weight 
of  the  flask  and  chlorine,  and  that  difference,  plus  the  weight 
of  the  air;  gives  the  weight  of  chlorine  occupying  the  given 
volume  under  the  second  set  of  conditions.  From  this  calculate 
the  weight  of  a  liter  of  chlorine. 

27.  Preparation  and  Properties  of  Chlorine.     Poison, 

a.  If  a  supply  of  chlorine  has  not  already  been  prepared  in  the 
previous  experiment  (26a.)  it  may  be  prepared  by  placing  about 
4  grams  of  manganese  dioxide  (Mn02)  in  a  generator  similiar 
to  that  used  for  hydrogen,  adding  through  the  thistle  tube  10 
c.c.  of  concentrated  hydrochloric  acid  diluted  with  10  c.c.  of 
water,  and  warming  gently.     The  delivery  tube  should  reach  to 
the  bottom  of  a  gas  collecting  bottle  (should  the  bottle  be  up- 
right or  inverted?)  which  should  be  covered  as  completely  as 
possible  with  a  glass  plate  and  the  bottle  should  be  replaced  as 
soon  as  the  greenish  color  of  the  chlorine  becomes  marked. 
Fill  three  bottles  in  this  manner,  cover  and  set  aside  and  im- 
mediately place  the  delivery  tube  of  the  generator  under  a 
dilute    potassium    or  sodium   hydroxide    solution;  force    the 
chlorine  out  of  the  generator  and  through  the  alkaline  solution 
(what  is  the  reaction  between  chlorine  and  the  alkali?)  by  pour- 
ing water  through  the  thistle  tube  until  the  generator  is  filled, 
then  empty  out  the  generator  before  proceeding. 

b.  Scatter  a  little  powdered  antimony  into  one  of  the  bottles 
of  chlorine     Write  the  reaction. 

c.  Place  strips  of  moist  and  of  dry  colored  calico  in  one  bottle 
and  account  for  the  difference  in  behavior. 

d.  Drop  moist  pieces  of  red  and  blue  litmus  paper  into 
the  other,  then  introduce  a  jet  of  burning  hydrogen  (from  a 
Kipp  generator)  until  the  appearance  of  the  flame  indicates 
that  the  chlorine  is  all  used  up,  then  again  drop  in  pieces  of 


42  THE  HALOGEN  FAMILY 

moist  litmus.  To  what  property  of  chlorine  is  the  effect  on 
litmus  due?  Why  must  the  litmus  be  moist?  What  is  the 
product  when  hydrogen  burns  in  chlorine,  and  what  is  its 
effect  on  litmus?  Would  you  expect  to  get  the  second  test 
with  litmus  if  any  free  chlorine  was  left  in  the  gaseous  mixture? 
Give  the  reason  for  your  answer. 

28.  Hydrochloric  Acid. 

a.  Review  the  experiment  in  which  hydrochloric  acid  was 
prepared  by  synthesis. 

b.  Calculate  the  amounts  of  sodium  chloride  and  sulfuric  acid 
required  by  the  equation  NaCl  +  H2S04  =  NaHS04  +  HCl  to 
give  enough  hydrochloric  acid  to  fill  two  gas-collecting  bottles, 
and  to  saturate  5  c.c.  of  water  at  20°  C.    Place  the  calculated 
amount  of  sodium  chloride  in  a  generating  flask  fitted  with 
a  thistle  tube  and  L-shaped  delivery  tube  and  add  through 
the  thistle  tube  a  little  more  than  the  calculated  amount  of 
concentrated  sulfuric  acid  (sp.  gr.  1.84),  diluted  with  one-half 

its  volume  of  water.  In  handling  concentrated 
sulfuric  acid  be  careful  not  to  let  it  touch  the 
skin,  as  it  causes  unpleasant  burns;  if  any 
does  get  on  the  skin  it  should  be  washed  off 
at  once  with  plenty  of  water.  In  diluting  sul- 
furic acid  always  pour  the  acid  into  the 
water.  Never  pour  water  into  concentrated  sul- 
furic acid. 

Place  the  generator  on  the  wire  gauze  and 
warm  gently,  and  with  the  delivery  tube  ex- 
tending to  the  bottom  of  the  gas  collecting 
bottle,  which  should  be  dry,  collect  two  bottles  of  the  gas 
by  displacement  of  air.  The  bottle  is  full  of  hydrochloric 
acid  when  moist  blue  litmus  paper  held  just  above  the  neck 
of  the  bottle  turns  red.  Cover  these  bottles  and  set  aside 
for  future  use.  Then  connect  the  delivery  tube  with  a  U 
tube  (storeroom)  as  shown  in  Fig.  13.  Place  5  c.c.  of  water 
in  the  U  tube  and  continue  the  generation  of  hydrochloric  acid 


HYDROCHLORIC  ACID  43 

till  the  water  is  saturated.     Save  the  generator  and  contents 
for  future  work. 

c.  Blow  across  the  mouth  of  one  of  the  bottles.     Explain  the 
formation  of  the  fumes.     Introduce  into  the  same  bottle  a 
glass   rod   wet   with   ammonium  hydroxide.     What   are   the 
fumes  in  this  case? 

d.  Invert  the  other  bottle  over  water  in  the  sink.    What  prop- 
erty of  hydrochloric  acid  does  this  illustrate?     If  any  gas  is 
left  in  the  bottle  after  shaking  about  with  the  water,  test  it 
to  determine  whether  it  is  hydrochloric  acid  or  air.     Explain 
the  test  and  results. 

Why  not  collect  the  hydrochloric  acid  by  displacement  of 
water?  How  would  you  determine  whether  a  solution  of 
hydrochloric  acid  is  saturated  or  not?  What  method  should 
be  used  to  determine  the  density  of  hydrochloric  acid  and 
why? 

e.  Place  1  c.c.  of  the  solution  from  the  U  tube  in  a  test-tube, 
and  add  a  granule  of  zinc.     What  is  the  reaction?    Pour  the 
solution  from  the  excess  of  the  zinc  and  evaporate  to  dryness 
in  an  evaporating  dish.     What  is  the  residue? 

/.  Add  portions  of  the  solution  to  a  small  amount  of  calcium 
oxide,  until  the  latter  is  entirely  dissolved.  Evaporate  this 
solution  to  dryness  in  an  evaporating  dish,  and  allow  it  to 
stand  over  night  in  the  dish  and  examine.  What  is  the 
reaction?  What  causes  the  change  on  standing? 

g.  Add  a  few  drops  of  silver  nitrate  (AgN03)  solution  to 
another  sample  of  the  hydrochloric  acid.  What  is  formed? 
Where  was  this  test  used  before?  For  what  radical  is  it  a 
test? 

h.  What  are  the  products  when  hydrochloric  acid  acts  on 
(a)  metals,  (b)  oxides,  (c)  hydroxides,  (d)  oxidizing  agents 
(Exp.  25)? 

i.  Pour  the  contents  of  the  generating  flask,  after  cooling,  into 
a  beaker  containing  about  10  c.c.  of  water  (or  if  the  contents 
of  the  flask  solidify  on  cooling  add  10  c.c.  of  water  to  the 
flask)  and  after  allowing  time  enough  for  the  solution  to  become 


44  THE  HALOGEN  FAMILY 

saturated  (fifteen  minutes  at  least)  pour  some  of  the  solution 
from  the  solid  residue  and  add  a  few  cubic  centimeters  of 
concentrated  hydrochloric  acid.  How  does  this  experiment 
prove  the  reaction  reversible?  What  means  may  be  used 
to  force  the  reversible  reaction  to  completion? 

29.  lonization. 

a.  Test  the  conductivity  of  each  of  the  following  by  placing 
about  10  c.c.  in  a  U  tube,  and  inserting  the  platinum  wires 
from  the  battery  into  each  arm.     If  a  current  flows;  the  am- 
meter which  is  connected  "in  series"  with  the  battery  will 
indicate  it  by  the  position  of  the  pointer.     Rinse  the  platinum 
wires  in  distilled  water  between  each  test. 

(a)  Distilled  water. 

(b)  Dilute  hydrochloric  acid. 

(c)  Dilute  potassium  hydroxide  (KOH)  solution. 

(d)  Dilute  sodium  chloride  (NaCl)  solution. 

(e)  Alcohol  and  water. 

(f)  Dilute  sugar  solution. 

Notice  the  phenomena  which  take  place  at  the  poles  when  a 
current  flows.  Test  to  determine  which  pole  gives  hydrogen  and 
which  oxygen  or  chlorine.  Which  pole  is  the  anode?  What 
happens  to  the  sodium  ion  when  it  loses  its  electrical  charge 
at  the  cathode,  when  electrolyzing  sodium  chloride  solution? 
Under  what  condition  could  metallic  sodium  be  prepared  by 
electrolysis  of  sodium  chloride?  The  first  products  of  the 
electrolysis,  in  this  case  sodium  and  chlorine,  are  called  the 
primary  products;  any  products  resulting  from  a  reaction  of 
the  primary  products  with  other  materials  present,  in  this 
case,  products  from  the  action  of  sodium  on  water  are  called 
secondary  products.  Write  equations  showing  the  primary  and 
the  secondary  products  of  each  electrolysis. 

b.  To  what  classes  of  chemical  compounds  do  conductors  of 
the  second  class  (those  which  conduct  in  solution,  and  at  the 
same  time  undergo  a  chemical  change)  belong?     How  does  the 
ionic  theory  explain  the  passage  of  the  current  through  the 


ROCK  SALT  45 

solution?  What  is  the  difference  between  the  chloride  ion  and 
free  chlorine  as  regards  (a)  color,  (b)  conduct  when  subjected 
to  the  influence  of  a  charged  plate  or  pole,  (c)  chemical  prop- 
erties. What  sort  of  reaction  is  especially  characteristic  of 
those  compounds  which  ionize  in  solution? 

30.  The  Purification  of  Rock  Salt,  (NaCl). 

Rock  salt,  as  mined,  is  of  course  mixed  with  more  or  less  of 
those  salts  which  are  in  solution  with  it  in  sea  water,  for  example 
sulfates  and  chlorides  of  calcium,  magnesium,  potassium,  etc. 
The  presence  of  sulfates  in  the  crude  salt  may  be  shown  by 
adding  a  few  drops  of  barium  chloride  (BaCl2)  solution  to  a 
dilute  solution  of  the  salt  acidified  with  C.  P.  hydrochloric  acid. 
A  white  precipitate  indicates  sulfates.  How  may  the  presence 
of  calcium  salts  be  shown?  (Exp.  21.) 

a.  Saturate  50  c.c.  of  water  with  common  salt.  From  the 
table  of  solubilities  the  amount  of  salt  required  may  be  deter- 
mined, and  a  little  more  than  the  theoretical  amount  should 
be  taken  to  allow  for  moisture  and  foreign  matter.  Filter  the 
solution  and  test  small  portions  of  it  for  sulfates  and  calcium. 
Divide  the  solution  remaining  into  two  equal  parts,  place  one 
part  in  an  evaporating  dish  and  evaporate  on  a  steam  bath 
until  about  two-thirds  of  the  water  has  evaporated,  allow  the 
salt  to  crystallize  by  cooling,  then  filter  out  the  crystals,  drain 
as  completely  as  possible,  wash  them  with  a  very  small  amount 
of  distilled  water,  drain,  label  No.  1,  and  set  away  to  dry.  Into 
the  other  half  of  the  solution  pass  hydrochloric  acid  gas,  gener- 
ated by  the  action  of  the  calculated  amount  of  concentrated 
sulfuric  acid  (sp.  gr.  1.84)  diluted  with  half  its  volume  of 
water,  on  10  grams  of  sodium  chloride.  The  generating  flask 
should  be  fitted  with  a  stopper  carrying  a  thistle  tube  reach- 
ing to  the  bottom  of  the  flask,  and  a  delivery  tube  dipping 
into  the  solution  of  sodium  chloride  to  be  purified.  Heat  the 
generating  flask  gently,  keeping  the  stream  of  gas  passing 
into  the  salt  solution  sufficiently  rapid  so  that  there  is  no 
tendency  for  the  solution  to  suck  back  into  the  generator,  and 


46  THE  HALOGEN  FAMILY 

continue  until  there  is  no  further  precipitation  in  the  salt 
solution.  Filter  out  the  precipitate  from  the  solution,  drain, 
wash  it  with  a  very  small  amount  of  distilled  water,  and 
drain,  finally  allowing  it  to  dry  completely  by  evaporation. 
Label  this  sample  No.  2.  What  is  this  precipitate?  Now 
dissolve  small  portions  of  each  sample  in  distilled  water 
and  test  as  before  for  sulfates  and  calcium.  Judge  the  rela- 
tive amounts  of  impurities  in  the  three  samples  tested,  and 
draw  conclusions  as  to  the  efficiency  of  the  two  methods  of 
purification.  Test  solutions  of  samples  1  and  2  with  red  and 
blue  litmus  paper.  What  has  become  of  the  hydrochloric 
acid  added  to  sample  2?  Is  the  precipitation  of  chlorides 
by  hydrochloric  acid  a  general  property?  Test  your  conclu- 
sion by  adding  1  c.c.  of  concentrated  hydrochloric  acid  to  1 
c.c.  of  saturated  solutions  (make  these  by  adding  an  excess 
of  the  solid  to  1  c.c.  of  water,  and  after  shaking  five  minutes 
pour  off  the  clear  solution  from  the  excess  of  the  solid)  of 
barium  chloride  (BaCl2)  and  potassium  chloride  (KC1). 

31.  Potassium  Hypochlorite. 

a.  Dissolve  5  grams  of  solid  potassium  hydroxide  (KOH)  in 
15  c.c.  of  distilled  water.  Cool  the  solution  and  place  about 
half  of  it  in  a  test-tube,  reserving  the  remainder  for  Exp.  32. 
Fit  up  a  small  generating  flask  with  a  delivery  tube  and  a  thistle 
tube  which  reaches  nearly  to  the  bottom  of  the  flask.  Place 
about  4  grams  of  finely  pulverized  manganese  dioxide  (Mn02) 
in  the  flask  and  add  10  c.c.  of  concentrated  hydrochloric  acid 
diluted  with  10  c.c.  of  water,  through  the  thistle  tube.  Warm 
the  flask  gently,  and  when  the  generation  of  chlorine  begins 
introduce  the  delivery  tube  into  the  test-tube  containing  the 
potassium  hydroxide  solution.  Surround  the  test-tube  with 
cold  water  and  stir  the  solution  while  chlorine  is  being  passed 
into  it.  Keep  a  lively  stream  of  chlorine  flowing  into  the  solu- 
tion for  ten  to  fifteen  minutes,  then  stop  the  generation  of 
chlorine  by  allowing  the  generator  to  cool  down,  remove  the 
test-tube  and  immerse  the  end  of  the  delivery  tube  in  water  or 
a  dilute  alkali  solution.  Do  not  allow  chlorine  to  escape  into  the 


POTASSIUM  CHLORATE  47 

air.     This  generator  may  be  used  again  in  Exp.  32,  but  it 
should  not  be  allowed  to  stand  over  night  in  the  desk. 

b.  Pour  the  contents  of  the  test-tube  into  an  evaporating 
dish  and  wet  a  piece  of  colored  cloth  in  it.  Is  the  cloth  bleached? 
Place  the  cloth  in  a  beaker  containing  some  dilute  sulfuric  acid. 
If  no  change  is  observed  wet  the  same  piece  of  cloth  with  the 
two  solutions  again.  Explain  what  happens.  Why  is  sulfuric 
acid  necessary?  Explain  why  this  treatment  will  bleach  cloth 
while  air  does  not.  Write  equations  for  the  reaction  in  a,  and 
for  the  action  of  sulfuric  acid.  Why  was  the  test-tube  cooled 
in  a  ?  Acidify  a  small  portion  of  the  solution  from  a  with  nitric 
acid,  then  add  a  few  drops  of  silver  nitrate  solution.  What  ion 
does  this  show  to  be  present? 

32.  Potassium  Chlorate. 

Using  the  same  generator  as  in  the  preceding  experiment  (31) 
pass  chlorine  into  the  remainder  of  the  potassium  hydroxide 
solution  but  do  not  cool  the  tube.  Continue  this  process  until 
the  solution  is  completely  saturated  with  chlorine,  then  dis- 
pose of  the  generator  and  contents  as  in  Exp.  27a.  How  is  it 
possible  to  tell  when  the  point  of  saturation  is  reached?  Allow 
the  mixture  to  cool.  What  happens?  Filter  and  test  the 
filtrate  by  adding  nitric  acid  and  silver  nitrate.  What  sub- 
stance is  present?  Dry  the  crystals  and  heat  in  a  dry  tube, 
testing  for  any  gas  that  might  be  evolved.  What  are  the 
crystals?  Which  is  most  soluble,  potassium  chloride  or 
potassium  chlorate?  Confirm  your  conclusion  by  reference 
to  the  table.  Write  all  reactions.  Explain  why  this  experi- 
ment and  the  preceding  do  not  yield  the  same  products. 

33.  The  Preparation  of  Potassium  Perchlorate. 

When  potassium  chlorate  is  heated  two  reactions  take  place 
simultaneously:  4  KC103  =  3KC104  +  KC1  and  2KC103  =  2KC1  + 
302.  The  greater  part  of  the  potassium  chlorate  will  go  through 
the  change  represented  by  the  first  equation  if  the  temperature 
is  held  just  above  the  melting-point  of  the  potassium  chlorate, 
so  that  the  evolution  of  oxygen  (represented  by  the  second 


48  THE  HALOGEN  FAMILY 

equation)  is  very  slow.  If  the  temperature  becomes  too  high 
the  evolution  of  oxygen  is  rapid,  the  amount  of  perchlorate 
formed  is  not  so  great,  and  it  may  itself  decompose  accord- 
ing to  the  equation,  KC104  =  KC1  +  202.  Thus,  in  preparing 
potassium  perchlorate  from  potassium  chlorate  the  potassium 
chlorate  must  be  heated  just  enough  to  keep  it  melted,  and 
evolving  a  slow  stream  of  oxygen. 

Place  about  8  grams  of  potassium  chlorate  (free  from  any 
combustible  matter,  why?)  in  the  hard  glass  test-tube,  and 
stopper  with  a  cork  or  rubber  stopper  carrying  a  delivery  tube 
leading  to  the  water  trough.  After  testing  the  apparatus  for 
air  tightness  melt  the  potassium  chlorate  and  heat  it  gently, 
collecting  the  oxygen  evolved  by  displacement  of  water. 
Take  care  that  the  melted  salt  does  not  come  in  contact  with 
the  stopper.  Why?  Regulate  the  heat  so  that  the  evolution 
of  oxygen  is  very  slow,  and  the  salt  is  kept  melted  throughout. 
When  500  c.c.  of  oxygen  have  been  evolved  withdraw  the  deliv- 
ery tube  and  allow  the  melted  mass  to  cool.  From  the  volume 
of  oxygen  evolved  calculate  the  weight  of  potassium  chlorate 
which  has  decomposed  according  to  the  second  equation. 
Calculate  the  amount  of  potassium  chloride  and  potassium  per- 
chlorate which  should  now  qe  in  the  reaction  mixture,  assuming 
that  all  the  potassium  chlorate  has  decomposed.  Look  up 
the  solubilities  of  the  two  salts  left  and  devise  a  method  for 
separating  these.  After  submitting  your  method  for  the  ap- 
proval of  an  instructor  carry  out  the  separation,  wash  the  po- 
tassium perchlorate  and  weigh  roughly.  What  per  cent  of  the 
theoretical  yield  is  obtained?  Should  the  method  of  sepa- 
ration used  give  a  pure  perchlorate?  Test  a  small  portion 
for  chlorides  and  save  the  rest  of  the  product  for  Exp.  34. 

34.  To  Test  the  Purity  of  the  Potassium  Perchlorate.    Quanti- 
tative. 

Place  from  0.2  to  0.5  gram  of  the  perfectly  dry  potassium 
perchlorate  obtained  in  Exp.  33  in  a  weighed  hard  glass  ignition 
tube,  weigh  accurately,  then  heat  until  the  evolution  of  oxygen 


BROMINE 


49 


is  complete,  and  weigh  the  tube  and  residue.  Calculate  the 
per  cent  of  oxygen  in  the  given  sample.  Compare  this  result 
with  that  obtained  in  Exp.  24  and  with  the  theoretical  compo- 
sition of  KC104.  Was  the  potassium  perchlorate  pure? 

35.  The  Preparation  of  Bromine. 

Caution:  Do  not  breathe  the  vapor  of  bromine.  Do  not 
get  liquid  bromine  on  the  skin,  as  it  causes  severe  burns. 

Calculate  the  quantities  of  manganese  dioxide  (Mn02), 
potassium  bromide  (KBr)  and  sulfuric  acid  necessary  to 
make  2  c.c.  of  liquid  bromine  (sp.  gr.  3.18).  Have  this  calcula- 


FIG.  14. 

tion  approved  by  an  instructor  before  proceeding.  The 
volume  of  concentrated  sulfuric  acid  required  (sp.  gr.  1.84) 
should  be  diluted  with  one-half  its  volume  of  water  by  pouring 
the  measured  amount  of  concentrated  acid  into  the  measured 
amount  of  water  in  a  beaker.  Never  pour  water  into  sulfuric 


50  THE  HALOGEN  FAMILY 

acid.  Grind  up  the  required  amount  of  potassium  bromide  in 
the  mortar  and  mix  with  the  manganese  dioxide,  then  place  the 
mixture  in  a  retort  (storeroom)  and  add  the  diluted  sulfuric 
acid  through  the  tubule.  Close  the  tubule  and  heat  gently, 
protecting  the  bottom  of  the  retort  by  a  wire  gauze.  Collect 
the  bromine  in  a  flask  into  which  the  neck  of  the  retort  extends 
as  far  as  possible,  and  which  is  kept  cool  by  being  placed  under 
a  stream  of  running  water,  while  the  escape  of  bromine  vapor 
into  the  air  is  prevented  by  folding  a  piece  of  wet  filter  paper 
over  the  mouth  of  the  receiving  flask  (Fig.  14).  When  the  evo- 
lution of  bromine  has  ceased  remove  the  flask  and  test  the 
gas  in  the  flask  with  pieces  of  paper  wet  with  starch  solution, 
and  with  starch  potassium  iodide  solution  (see  Exp.  25) .  (Then 
proceed  at  once  to  Exp.  36.  Do  not  allow  the  bromine  to  stand 
over  night.)  To  what  is  the  difference  in  color  in  the  two  tests 
due?  Refer  to  the  starch  potassium  iodide  test  for  chlorine. 
Would  you  expect  to  find  the  same  succession  of  colors  in  this 
case?  Write  equations  for  the  reactions  involved.  Could 
bromine  be  prepared  by  any  of  the  methods  used  for  chlorine? 
If  potassium  iodide  was  present  as  an  impurity  in  potassium 
bromide  how  could  the  potassium  bromide  be  most  readily 
purified?  Is  there  any  connection  between  the  low  boiling- 
point  of  bromine  and  the  fact  that  the  air  above  the  liquid 
bromine  is  strongly  colored  with  bromine  vapor  at  room 
temperature? 

36.  Potassium  Bromate. 

a.  Cover  the  bromine  collected  in  the  preceding  experiment 
with  2  to  4  c.c.  of  water  and  add  drop  by  drop  a  strong  solution 
of  potassium  hydroxide  (KOH,  1:1)  with  shaking,  until  the 
liquid  bromine  has  entirely  dissolved,  and  the  color  of  the  solu- 
tion is  only  faintly  yellow.  What  reaction  is  taking  place? 
(Note:  If  a  brown  precipitate  of  manganese  dioxide,  carried  over 
mechanically  by  the  bromine,  appears  in  the  solution  on  the 
addition  of  potassium  hydroxide  it  must  be  filtered  out  before 
the  color  of  the  bromine  in  solution  is  entirely  removed  by  the 


POTASSIUM  BROMATE  51 

potassium  hydroxide.)  Heat  this  solution  to  boiling,  (what 
further  reaction  occurs?),  and  continue  boiling  until  crystals 
begin  to  separate  out,  then  cool  under  running  water,  and  filter 
after  the  solution  has  come  to  room  temperature.  Rinse  out 
the  flask  and  the  precipitate  on  the  filter  paper  with  2  c.c.  water 
and  save  the  filtrate.  Allow  the  crystals  to  drain  as  com- 
pletely as  possible  and  then  in  order  to  purify  them  wash 
into  an  evaporating  dish  with  a  fine  stream  of  distilled 
water  from  the  wash  bottle.  Heat  till  the  crystals  are 
entirely  dissolved,  then  allow  the  solution  to  crystallize 
spontaneously  by  cooling  and  evaporation.  Remove  the 
crystals  from  the  solution  before  all  the  water  has  evapo- 
rated and  dry  them  on  filter  paper,  then  place  in  a  weighed 
(quantitative)  test-tube  and  label  in  such  a  way  that  the  label 
can  be  readily  removed  before  a  second  weighing.  Boil  the 
filtrate  from  the  first  crystallization  (in  an  evaporating  dish) 
until  it  begins  to  crystallize,  then  allow  to  cool.  Remove  the 
crystals,  dry  them  on  filter  paper,  and  place  them  in  another 
weighed  test-tube,  labeling  as  before.  These  may  be  best 
labeled  according  to  solubility.  Which  crystals  represent  the 
least  soluble,  and  which  the  most  soluble  portion?  Write 
the  reaction  involved  when  bromine  acts  on  potassium  hydrox- 
ide under  the  given  conditions,  and  after  looking  up  the  solu- 
bilities of  the  products  determine  which  salt  you  will  expect 
to  find  in  the  least,  and  which  in  the  most  soluble  portion  of 
your  product.  Weigh  the  test-tubes  with  the  salt,  and  by 
subtracting  the  weight  of  the  empty  tubes  get  the  weights  of 
the  salts. 

b.  Devise  some  method  of  testing  these  two  salts  to  determine 
whether  your  conclusions  regarding  their  identity  are  correct. 
Get  the  approval  of  an  instructor  before  proceeding  to  apply  the 
tests  you  devise.  Compare  the  reaction  between  bromine  and 
potassium  hydroxide  with  that  used  in  Exp.  32  between 
chlorine  and  potassium  hydroxide.  Calculate  from  the  original 
amount  of  KBr  taken  in  Exp.  35,  the  theoretical  yields  of 
KBr03  and  KBr.  What  per  cent  of  the  theoretical  yield  is 


52  THE  HALOGEN  FAMILY 

the  amount  obtained;  or  in  other  words,  what  is  your  per- 
centage yield?  At  what  stages  of  the  process  has  there  been  an 
opportunity  for  loss  of  material? 

37.  Iodine. 

Mix  together  a  small  crystal  of  potassium  iodide,  and  a 
little  manganese  dioxide.  Place  the  mixture  in  a  test-tube 
and  add  a  few  drops  of  concentrated  sulfuric  acid,  then  heat 
gently.  Note  the  color  of  the  gas  evolved.  Is  it  lighter  or 
heavier  than  air?  What  observation  leads  to  this  conclusion? 
Hold  a  paper  wet  with  starch  solution  in  the  gas.  Explain. 
Why  was  potassium  iodide  used  with  the  starch  solution  in  the 
test  for  chlorine  and  bromine  and  not  in  this  test?  Compare 
this  method  for  the  preparation  of  iodine  with  those  for  chlorine 
and  bromine.  Why  would  it  not  be  safe  to  use  potassium 
chlorate  as  the  oxidizing  agent  on  a  mixture  of  potassium 
iodide  and  concentrated  sulfuric  acid  (see  test-book  under 
chlorine  dioxide).  What  would  be  the  reaction  between  KC103 
and  HI? 

38.  Hydrofluoric  Acid  and  its  Etching  Action  on  Glass. 

Prepare  hydrofluoric  acid  by  placing  2  grams  of  calcium 
fluoride  (CaF2)  in  a  lead  dish,  or  evaporating  dish  which  has 
been  coated  with  paraffin,  and  mixing  with  enough  concentrated 
sulfuric  acid  to  make  a  paste.  Cover  the  dish  with  a  glass 
plate  which  has  been  protected  by  a  coating  of  paraffin  except 
in  the  places  which  are  to  be  etched.  Allow  the  dish  and 
cover  to  stand  for  about  an  hour,  then  remove  the  paraffin 
from  the  cover  glass  and  examine.  Write  the  reactions  and 
account  for  the  disappearance  of  part  of  the  glass.  Test  the 
gas  in  the  dish  by  breathing  across  it.  What  are  the  fumes? 
Hold  a  rod  wet  with  ammonium  hydroxide  (NH4OH)  above 
the  dish.  What  are  the  fumes?  Is  this  a  distinctive  test  for 
hydrofluoric  acid?  For  what  other  acid  has  it  been  used? 
What  one  property  is  essential  that  an  acid  should  give  this 
test? 


EQUILIBRIUM  53 

39.  Equilibrium. 

a.  Review  the  preparation  of  hydrochloric  acid  and  show  how 
the  reaction  was  proved  to  be  reversible.     What  property  of 
the  hydrochloric  acid  made  it  possible  to  get  complete  con- 
version of  the  sodium  chloride  into  hydrochloric  acid,  even 
though  the  reaction  was  reversible? 

The  reaction,  NaH2P04  +  HC1  «  NaCl  +  H3P04,  is  revers- 
sible,  and  if  at  any  fixed  temperature  solutions  of  primary 
sodium  phosphate  and  hydrochloric  acid,  or  of  sodium  chloride 
and  phosphoric  acid  of  equivalent  concentrations  are  mixed, 
the  reaction  mixture  will  consist  of  a  solution  of  the  two 
original  materials  and  the  two  resulting  products  in  exactly 
the  same  proportions,  regardless  of  which  pair  of  materials 
was  originally  used. 

b.  To  show  this  experimentally  make  an  indicator  by  mixing 
six  volumes  of  potassium  iodide  solution  with  one  volume  of 
potassium  bromate  solution,  and  add  a  few  drops  of  this  indi- 
cator to  small  portions  of  normal  hydrochloric  and  phosphoric 
acid  solutions,  noting  the  effect  of  the  indicator  on  the  two 
acids.    Now  prepare  four  tubes,  in  one  placing  equal  quantities 
of  normal  solutions  of  primary  sodium  phosphate  and  hydro- 
chloric acid,  in  another  the  same  quantities  of  normal  solutions 
of  sodium  chloride  and  phosphoric  acid,  in  the  third  the  same 
quantities  of  normal  hydrochloric  acid  and  water,  in  the  fourth 
the  same  quantities  of  normal  phosphoric  acid  and  water,  and 
label  each  tube.     Shake  each  thoroughly  then  add  an  equal 
number  of  drops1  of  the  indicator  to  each  (the  two  diluted 
acid  solutions  are  to  be  used  as  standards  to  show  the  color 
given  by  hydrochloric  and  phosphoric  acids  of  the  same  con- 
centration as  in  the  first  two  tubes)  .     From  the  color  reactions 
(these  must  be  noted  at  once,  as  the  color  deepens  on  standing) 
draw  conclusions  as  to  the  composition  of  the  solution  in  each 
of  the  first  two  tubes. 


student  should  provide  himself  with  half  a  dozen  pieces  of  glass 
tubing,  drawn  to  a  jet  at  one  end  and  kept  clean  and  dry  to  use  as  drop- 
ping tubes. 


54  THE  HALOGEN  FAMILY 

c.  Add  twice  normal  phosphoric  acid  to  an  equal  volume  of 
normal  sodium  chloride  and  note  the  color  given  by  the  same 
number  of  drops  of  the  indicator  as  before.  Compare  with  the 
tube  made  from  normal  phosphoric  acid  and  sodium  chloride. 
What  effect  does  increased  concentration  of  one  of  the  reacting 
substances  have  on  the  point  of  equilibrium? 


CHAPTER  VII 
SULFUR 

40.  Physical  Properties  of  Sulfur. 

What  is  a  polymorphous  substance?  Do  the  polymorphic 
forms  of  a  given  substance  differ  most  with  respect  to  physical 
or  chemical  properties?  Illustrate.  Place  about  5  grams 
of  sulfur  (sample  A)  in  a  short,  wide,  dry  test-tube  and  heat 
gently  till  it  just  melts  (use  special  care  not  to  heat  the  sulfur 
much  above  its  melting-point),  then  pour  a  small  portion  of 
the  melted  sulfur  into  a  beaker  containing  cold  water  (sample 
B) .  Continue  the  heating  of  the  rest  of  the  sulfur  in  the  test- 
tube  until  it  begins  to  boil,  noting  the  changes  in  appearance 
and  viscosity  as  the  temperature  rises.  Pour  the  boiling 
sulfur  into  another  beaker  of  cold  water  (sample  C).  Com- 
pare these  three  samples  of  sulfur  (A,  B,  C)  as  to  appearance, 
texture  and  solubility  in  carbon  disulfide.  To  study  this 
last  property  place  a  small  portion  of  each  sample  in  separate 
test-tubes  and  add  1-2  c.c.  of  carbon  disulfide  (CS2),  then 
shake,  taking  care  that  the  test-tubes  are  at  a  safe  distance 
from  all  flame,  as  carbon  disulfide  is  very  inflammable,  and 
its  vapor  mixed  with  air  is  explosive.  Allow  small  portions  of 
these  solutions  to  evaporate  spontaneously.  Did  any  sulfur 
dissolve?  Record  the  results,  then  pour  the  contents  of  the 
test-tubes  down  the  sink,  following  with  plenty  of  water.  Let 
portions  of  samples  B  and  C  stand  in  the  desk  over  night  and 
examine  again  in  the  same  way.  Explain  the  differences  in 
the  behavior  of  the  three  samples.  Retain  the  test-tube  in 
which  the  sulfur  was  melted  for  use  in  the  next  experiment 
(41  6.) 

55 


56  SULFUR 

41.  Chemical  Properties  of  Sulfur  and  Hydrogen  Sulfide. 

a.  Mix  1  gram  powdered  zinc  and  0.5  gram  sulfur,  place  on  a 
crucible  cover  and  heat.     Caution.     What  is  the  reaction? 

b.  Mix  2  grams  of  iron  filings  and  1  gram  of  sulfur,  place  in 
the  test-tube  used  in  Exp.  40  and  heat.     What  is  the  reaction? 

c.  What  reactions  of  chlorine  and  oxygen  are  analogous  to 
these  for  sulfur?    From  the  position  of  sulfur  in  the  periodic 
table  would  you  expect  it  to  be  like  these  two  elements? 
Examine  these  two  products  by  placing  samples  of  each  (obtain 
the  iron  compound  by  breaking  the  tube)  in  separate  test-tubes 
and  adding  a  few  c.c.  of  dilute  hydrochloric  acid.     Note  the  odor 
of  the  gas  evolved  (cautiously).     Test  it  with  paper  moistened 
with  lead  acetate  solution.     Write  the  reactions  involved. 

d.  Introduce  a  small  portion  of  burning  sulfur  on  the  iron 
spoon  into  a  gas-collecting  bottle,  and  cover  the  bottle  as  com- 
pletely as  possible  with  a  glass  plate.     What  is  the  reaction? 
When  the  reaction  has  ceased  place  10  c.c.  of  water  in  the  bottle 
and  shake.     What  is  formed?    Pour  this  solution  into  a  test- 
tube  and  pass  in  hydrogen  sulfide,  which  may  be  generated  by 
placing  some  of  the  iron  or  zinc  sulfide  in  a  test-tube,  adding 
dilute  hydrochloric  acid  and  closing  the  tube  with  a  stopper 
carrying  a  delivery  tube  which  can  be  inserted  into  the  solution. 
Note  the  appearance  and  odor  of  the  solution  after  the  reaction 
has  taken  place.     Write  the  equation. 

e.  This  is  an  example  of  a  colloidal  solution,  and  is  called 
"milk  of  sulfur."     Two  of  the  general  properties  of  colloidal 
solutions  may  be  easily  illustrated   by  pouring  the   solution 
through  a  filter  paper,  showing  the  fine  state  of  division  of  the 
sulfur,  then  adding  a  few  crystals  of  sodium  chloride,  shaking 
and  filtering  again.     The  electrolyte  has  coagulated  the  colloid. 

/.  Pass  some  hydrogen  sulfide,  generated  in  the  same  way, 
into  10  c.c.  of  water,  stopper  and  allow  the  solution  to  stand 
over  night.  What  is  formed?  What  is  the  oxidizing  agent  in 
this  case?  These  reactions  have  been  suggested  to  explain  the 
occurrence  of  free  sulfur  in  volcanic  regions. 

g.  Pass  hydrogen  sulfide  from  a  Kipp  generator  or  from  the 


SULFUR  DIOXIDE  57 

laboratory  supply  into  a  solution  of  a  zinc  salt,  for  example 
zinc  chloride.     Is  the  reaction  reversible? 

42.  Sulfur  Dioxide. 

a.  If  fresh  solutions  of  ferric  chloride  (FeCl3)  and  potassium 
ferricyanide  (K3Fe(CN)6)  are  mixed,  and  the  brown  solution 
so  obtained  is  taken  up  on  filter  paper,  it  may  be  used  as  an 
indicator  for  sulfur  dioxide  (S02),  as  the  color  changes  to  a 
bluish-green  when  exposed  to  the  gas.     The  color  changes  in  the 
same  way  if  the  paper  is  exposed  to  air  and  light  for  any  length 
of  time,  so  it  must  be  prepared  and  used  for  the  test  at  once. 
If  this  is  done  there  is  no  difficulty  in  recognizing  appreciable 
amounts  of  sulfur  dioxide. 

b.  Heat  2  c.c.  of  concentrated  sulfuric  acid  in  contact  with 
copper  turnings  in  a  test-tube  until  there  is  a  spontaneous 
evolution   of  gas.     Test   the  escaping  gas   with   the  freshly 
prepared  paper.     What  is  formed?     Write  the  reaction.     Does 
dilute  sulfuric  acid  act  on  copper?     (See  Exp.  13.)     Allow  the 
tube  and  contents  to  cool,  then  pour  the  contents  in  a  beaker 
containing  about  50  c.c.  of  water.     Allow  the  precipitate  to 
settle,  and  explain  the  color  of  the  solution. 

c.  Add  concentrated  sulfuric  acid  to  a  small  crystal  of  potas- 
sium bromide  (KBr)  and  heat  gently,  testing  the  escaping  gas  as 
before.     Explain  the  color  of  the  evolved  gas  in  this  case. 
Write  the  reactions  taking  place. 

d.  What  name  would  be  given  to  any  process  by  which  S02  is 
obtained  from  H2S04?     In  these  two  experiments  what  are  the 
reducing  agents?     What  is  the  oxidizing  agent? 

e.  Add  some  dilute  hydrochloric  acid  to  a  few  sodium  sulfite 
(Na2S03)  or  sodium  acid  sulfite  (NaHS03)  crystals  in  a  test- 
tube,  warm  and  test  for  sulfur  dioxide.     Write  the  reaction. 
Fit  this  tube  with  a  cork  and  delivery  tube,  and  pass  the  gas 
evolved  (use  care  that  none  of  the  liquid  is  carried  over)  into 
a  test-tube  containing  9  c.c.  of  water  and  1  c.c.  ammonium 
hydroxide    (NH4OH)   until  the  solution  becomes  saturated. 
Now   pass  the  gas  into  another  test-tube   containing  1  c.c. 


58  SULFUR 

potassium  permanganate  (KMn04)  solution  and  9  c.c.  water. 
Note  the  change  in  appearance.  (What  other  reagent  has 
the  same  effect  on  KMn04?) 

Compare  the  two  solutions  by  adding  to  each  a  little  barium 
chloride  (BaCl2)  solution,  shake  and  note  the  effect,  then  allow 
to  settle,  decant  the  clear  liquid  and  add  dilute  C.P.  hydro- 
chloric acid  to  the  precipitate  in  each  tube.  How  do  you 
account  for  the  difference  in  behavior?  Write  equations  for 
all  the  reactions  involved. 

43.  Preparation  and  Properties  of  Sodium  Thiosulfate. 

Dissolve  10  grams  of  anhydrous  sodium  carbonate  in  50  c.c. 
of  water,  warming  if  necessary,  or  if  only  the  hydrated  salt 
(Na2C03-10  H20)  is  available  dissolve  a  quantity  of  it  equivalent 
to  10  grams  of  Na2C03  in  a  correspondingly  smaller  amount  of 
water.  Divide  this  solution  into  two  equal  parts,  and  saturate 
one  part  with  sulfur  dioxide,  generated  by  heating  10  grams 
of  copper  turnings  with  20  c.c.  concentrated  sulfuric  acid  in  a 
generating  flask  provided  with  a  thistle  tube  and  delivery  tube. 
Be  particularly  careful  to  see  that  the  solution  of  the  carbonate  is 
not  at  any  time  sucked  back  into  the  generating  flask.  Add  to 
the  solution  prepared  in  this  way  (what  is  it?)  the  other  part 
of  the  sodium  carbonate  solution,  and  heat  to  boiling.  What 
two  reactions  have  occurred  thus  far?  Write  equations  for 
each.  After  the  effervescence  has  ceased  (to  what  is  it  due  ?) 
add  4  grams  of  sulfur,  boil  for  20  minutes  keeping  the  volume 
of  the  solution  constant,  filter  off  the  excess  of  sulfur,  evapo- 
rate the  solution  over  a  water  bath  until  crystallization  com- 
mences, then  allow  it  to  cool.  Filter  off  the  crystals  and  evapo- 
rate the  filtrate  to  crystallization  a  second  time,  then  cool  and 
combine  the  two  crops  of  crystals.  Weigh  the  product  (after 
drying)  on  the  rough  laboratory  balance,  and  calculate  the 
percentage  yield.  Dissolve  some  of  the  crystals  in  water, 
and  test  the  solution  with  blue  and  red  litmus  paper. 
Hydrolysis  of  the  salt  into  a  base  and  acid  has  taken  place. 
According  to  the  test,  which  (base  or  acid)  is  more  highly 


SULFURIC  ACID  59 

ionized?  Add  a  few  drops  of  hydrochloric  acid  to  the  solution 
and  warm.  The  precipitate  is  milk  of  sulfur,  where  have  you 
seen  this  before?  Write  the  equation  for  the  reaction.  Is 
thiosulfuric  acid  a  stable  compound?  Add  a  solution  of  so- 
dium thiosulfate  to  some  freshly  prepared  silver  chloride  and  to 
a  small  crystal  of  iodine.  Write  the  reactions.  Of  what  tech- 
nical use  is  sodium  thiosulfate? 

44.  Sulfuric  Acid. 

a.  Study  the  properties  of  sulfuric  acid  by  the  following  tests : 
Dip  a  strip  of  filter  paper  into  dilute  sulfuric  acid  and  dry 
on  the  radiator,  then  examine.     Add  a  few  drops  of  concen- 
trated sulfuric  acid  to  a  little  sugar  in  an  evaporating  dish. 
Compare  with  the  effect  of  heating  sugar  on  a  crucible  lid. 
Add  1  c.c.  of  concentrated  sulfuric  acid  to  1  c.c.  of  water  in 
a  test-tube,  noting  the  temperature  before  and  after  mixing. 
Does  the  last  experiment  throw  any  light  upon  the  nature  of 
the  change  in  the  first  two?     Explain. 

b.  Add  barium  chloride  (BaCl2)   solution  to  dilute  sulfuric 
acid.     What  is  the  reaction?     Allow  the  precipitate  to  settle, 
pour  off  the  clear  liquid  and  add  a  few  cubic  centimeters  of  con- 
centrated C.  P.  hydrochloric  acid  to  the  precipitate.     Is  the 
above  reaction  appreciably  reversible  under  these  conditions? 

45.  Neutralization.     Determination   of  the   Concentration   of   a 
Solution  of  Acid  or  Base.     Quantitative. 

The  neutralization  of  an  acid  by  a  base  depends  on  the  union 
of  the  hydrogen  ions  of  the  acid  with  the  hydroxyl  ions  of  the 
base  to  form  unionized  water.  If  an  indicator,  litmus,  is 
present  in  a  solution  of  an  acid  it  will  give  the  solution  a  red 
color,  due  to  the  H+  ions.  If  a  base,  as  KOH,  is  added  to  that 
acid  solution,  the  H+  ions  are  gradually  removed  by  the  OH" 
ions  of  the  base,  until  finally  there  are  none  left  in  the  solution, 
and  the  addition  of  the  next  drop  of  the  solution  of  the  base 
causes  the  litmus  to  turn  blue,  because  of  the  free  OH~  ions 
in  the  solution.  Such  a  solution,  the  addition  of  one  drop  of 


60  SULFUR 

base  or  acid  to  which  causes  it  to  turn  litmus  blue  or  red,  is 
said  to  be  neutral,  and  is  free  from  both  H+  and  OH~  ions 
except  to  the  extent  to  which  they  occur  in  water,  where  they 
are  equal  in  number.  What  ions  are  present  in  a  solution  of 
H2S04  which  has  been  neutralized  by  KOH?  Write  the  ionic 
equation  for  the  reaction. 

The  concentration  of  an  acid  solution  may  be  determined 
by  neutralizing  a  measured  volume  of  the  solution,  evaporating 
to  dryness  and  weighing  the  salt  formed  (it  must  be  a  non- 
volatile, stable  salt,  why?)  and  calculating  from  the  weight  of 
the  salt,  the  weight  of  the  acid  originally  in  the  solution. 

a.  Make  a  dilute  solution  of  sulfuric  acid  by  pouring  3  c.c. 
of  concentrated  sulfuric  acid  into  100  c.c.  of  water  and  mix 
thoroughly  by  shaking.     Place  in  a  flask  and  stopper  the  flask, 
so   that   the  solution  will   not   change  in   concentration   by 
evaporation.     Dissolve    about    6    grams    of   solid    potassium 
hydroxide  (KOH)  (use  care  in  handling  this,  as  it  is  a  caustic 
and  burns  the  skin)  in  100  c.c.  of  water,  shake  thoroughly 
and  stopper.     Label  the  flasks. 

Save  these  solutions  as  they  must  be  used  again  in  Exp.  46. 

b.  Rinse  out  and  fill  a  burette  (storeroom)  with  the  prepared 
potassium  hydroxide  solution  drawing  enough  solution  out 
through  the  nozzle  to  insure  the  removal  of  all  the  air  from  the 
fine  tubing,  and  bring  the  solution  in  the  burette  to  the  zero 
mark.     Now  make  a  rough  determination  of  the  amount  of 
potassium  hydroxide  solution  required  to  neutralize  10  c.c.  of 
the  sulfuric  acid  solution  as  follows:  Measure  out  10  c.c.  of 
the  sulfuric  acid  solution  with  the  pipette,  and  place  it  in  a 
beaker,  add  2  drops  of  phenolphthalein  (which  is  an  indicator 
like  litmus,  except  that  it  is  colorless  in  acid  and  red  in  alkaline 
solutions)    and    then    while    stirring    continuously,    add    the 
potassium  hydroxide  solution  from  the  burette,  1  c.c.  at  a  time, 
until  the  solution  in  the  beaker  becomes  permanently  red,  i.e., 
until    it    is    alkaline.     Read    the    burette.     The    volume    of 
potassium  hydroxide  solution  used  is  not  more  than  1  c.c. 
greater  than  the  volume  required  for  exact  neutralization  of 


NEUTRALIZATION  61 

10  c.c.  of  the  sulfuric  acid  solution  if  these  directions  have 
been  carefully  followed.  Now  make  the  exact,  final  deter- 
mination as  follows: 

c.  Clean  and  weigh  (quantitative)  an  evaporating  dish  and 
cover,    then   place  in  it  exactly  10  c.c.   of  the  sulfuric  acid 
solution,  measured  with  the  pipette,  and  weigh  again.     From 
the  weight  of  the  solution  and  its  volume  calculate  its  specific 
gravity.     Add  a  couple  of  drops  of  phenolphthalein  and  after 
reading    the    burette    neutralize    the    acid    with    potassium 
hydroxide  from  the  burette,  adding  at  once  1  c.c.  less  than  the 
amount  required  in  the  rough  neutralization,  then  to  determine 
the  fraction  of  a  cubic  centimeter  required  to  just  neutralize 
the  acid  add  the  potassium  hydroxide  drop  by  drop,  stirring 
constantly  with  a  glass  rod,  and  stop  the  addition  the  moment 
the  solution  turns  permanently  red.     The  color  is  best  seen 
against  a  white  background. 

When  the  solution  is  just  neutral  read  the  burette  again,  and 
determine  the  volume  of  potassium  hydroxide  solution  used. 
Evaporate  the  neutral  solution  to  dryness  on  a  steam  bath,  or 
by  setting  it  above  a  beaker  of  boiling  water,  and  when  dry 
heat  the  covered  dish  gently  with  the  direct  flame,  cool  and 
weigh.  Heat  the  dish  and  contents  again,  and  weigh,  until 
two  successive  weighings  give  the  same  result.  What  is  left 
in  the  dish? 

d.  Define  per  cent  composition  and  normal  solution.    From  the 
weight   of   the   salt  formed  calculate  the  weights  of  sulfuric 
acid  and  potassium  hydroxide  used.     Calculate  the  normal 
strength    of    the    sulfuric     acid    and    potassium    hydroxide 
solutions.     How  many  milligrams  of  sodium  hydroxide  (NaOH) 
will  1  c.c.  of  the  sulfuric  acid  neutralize?     How  many  milli- 
grams of  ammonium  hydroxide  (NH4OH)?     How  many  milli- 
grams  of  nitric   acid   (HN03)   will   1   c.c.   of  the   potassium 
hydroxide   neutralize?     Of  hydrochloric   acid?     Of  sulfurous 
acid    (H2S03)?     Calculate   the   per   cent   composition   of  the 
sulfuric  acid  solution.     What  further  information  would  be 
necessary  for  you  to  calculate  the  per  cent  composition  of  the 


62  SULFUR 

potassium  hydroxide  solution?  Why  could  not  the  concen- 
tration of  these  solutions  be  determined  by  evaporating  off 
the  water  and  weighing  the  residue,  as  in  Exp.  18? 

46.  Potassium  Acid  Sulfate. 

Place  10  c.c.  of  the  dilute  sulfuric  acid  prepared  in  Exp.  45 
(measured  with  the  pipette)  in  a  clean  evaporating  dish,  then 
add,  from  the  burette,  just  half  as  much  potassium  hydroxide 
solution  as  was  found  necessary  for  neutralization,  and  evapo- 
rate the  solution  to  dryness  on  the  steam  bath.  What  is 
formed?  Dissolve  a  few  of  the  crystals  in  water  and  test  the 
solution  with  red  and  blue  litmus  paper.  Explain  the  effect. 
Dip  a  filter  paper  into  the  solution,  and  allow  it  to  dry  out  on 
the  radiator.  Is  sulfuric  acid  present  in  the  solution?  If 
sulfuric  acid  is  not  found,  how  can  the  effect  on  litmus  be 
explained?  Add  some  of  the  solution  to  some  solid  sodium 
carbonate  in  a  test-tube.  Explain.  Add  barium  chloride 
(Bad 2)  to  the  solution,  then  hydrochloric  acid.  Is  the  sulfate 
ion  present? 


CHAPTER  VIII 

NITROGEN 

47.  Preparation  and  Properties  of  Nitrogen  and  Ammonia. 

a.  What   are  the  principal  constituents   of  the  air?     Are 
they   present    as    a   mixture    or   in    chemical    combination? 
How  may  oxygen  be  removed  from  the  air? 

b.  Float  a  crucible  on  the  trough  full  of  water,  place  in  it  about 
0.5  gram  red  phosphorus  and  light  it  by  touching  it  with  the 
hot  end  of  a  file  or  iron  spoon.     Quickly  invert  a  gas-collecting 
bottle  over  the  dish,  bringing  the  mouth  of  the  bottle  below  the 
surface  of  the  water,  and  hold  in  that  position  till  the  phosphorus 
has  ceased  burning,  and  the  water  inside  the  bottle  has  come  to  a 
fixed  level.     Account  for  the  rise  of  the  water  in  the  bottle. 
Place  a  glass  cover  over  the  mouth  of  the  bottle  while  it  is  still 
under  water,  and  withdraw.     Test  the  gas  in  the  bottle  by  in- 
serting into  it  a  burning  splinter.     Does  it  support  combustion? 
Does  it  burn?     Insert  into  it  a  piece  of  burning  magnesium 
ribbon.     Does  it  support  combustion  in  this  case?     What  is 
formed? 

c.  Place  about  0.5  gram  of  magnesium  powder  in  a  crucible, 
cover  and  heat  while  covered  until  all  glowing  has  ceased. 
The  crucible  contains  a  mixture  of  what  two  compounds? 
Place  the  contents  of  the  crucible  in  a  test-tube,  moisten  with 
water,  and  warm  gently.     Test  the  escaping  gas  with  moistened 
red  litmus  paper.     What  is  it?    Write  the  reactions. 

d.  To  1  c.c.  of  ammonium  chloride  (NH4C1)  solution  in  a  test- 
tube  add  a  few  drops  of  potassium  or  sodium  hydroxide  solution, 
and  warm,  testing  the  gas  evolved  as  before.     What  is  it? 
Write  the  reaction.     Hold  a  rod  wet  with  concentrated  hydro- 
chloric acid  at  the  mouth  of  the  test-tube.     Explain  the  fumes 
formed.     Devise  and  apply  a  test  to  determine  whether  the 
gas  is  lighter  or  heavier  than  air. 

63 


64  NITROGEN 

48.  The  Weight  of  a  Liter  of  Ammonia.     Quantitative. 

Generate  ammonia  by  warming  together  in  flask  fitted  with  a 
delivery  tube,  and  proved  to  be  air  tight,  a  mixture  of  equal 
weights  of  ammonium  chloride  and  powdered  quick  lime  (CaO). 
Collect  three  test-tubes  of  the  gas  by  air  displacement  and  test 
its  solubility  in  water,  dilute  alkalis  (dilute  sodium  hydroxide 
or  potassium  hydroxide)  and  dilute  acids.  From  this  test 
determine  which  will  make  the  most  suitable  liquid  for  absorbing 
the  gas,  then  proceed  with  the  determination  of  the  weight  of  a 
liter,  according  to  the  method  used  for  chlorine.  Could  con- 
centrated sulfuric  acid  be  used  to  dry  this  gas?  Why?  If  the 
gas  is  prepared  as  directed,  from  dry  ammonium  chloride  and 
quicklime  it  will  not  be  necessary  to  dry  it,  but  it  may  be  passed 
directly  into  the  weighed  flask.  Should  the  gas  be  collected 
with  the  flask  held  upright  or  inverted?  The  flask  is  sufficiently 
filled  when  a  strip  of  moistened  red  litmus  paper  held  at  the 
mouth  turns  blue. 

49.  Ammonium  Hydroxide. 

a.  Attach  a  U  tube  containing  water  to  the  ammonia  generator 
used  in  the  previous  experiment,  similar  to  the  arrangement  used 
for  making  a  solution  of  hydrochloric  acid  (Fig.  13)  and  pass  in 
ammonia  until  the  solution  is  saturated.     Note  the  effect  on 
the  temperature  of  the  solution.     Test  this  solution  with  litmus, 
then  boil  a  small  portion  in  a  beaker  for  5  minutes  and  test  it 
again.     Explain.     Neutralize  another  portion  of  the  solution 
with  dilute  hydrochloric  acid  and  evaporate  the  neutral  solution 
to  dryness  on  the  water  bath,  then  heat  the  dry  residue  with  the 
direct  flame.     Repeat  this  procedure  only  neturalize  with  dilute 
sulfuric  acid  instead  of  hydrochloric.     Account  for  the  difference 
in  behavior  on  heating  the  dry  salts.     Is  there  any  relation  be- 
tween the  volatility  of  the  acid  and  of  its  ammonium  salt? 
Allow  a  small  sample  of  ammonium  carbonate  to  stand  on  a 
watch  glass  over  one  laboratory  period,  and  account  for  the 
change. 

b.  Add  ammonium  hydroxide  solution  one  drop  at  a  time  to 


NITRIC  ACID  65 

freshly  precipitated  silver  chloride  and  to  copper  sulfate  solution. 
These  reactions  are  probably  due  to  the  presence  of  ammonia, 
(NH3)  in  the  solution,  and  the  soluble  compounds  formed  are 
AgCl-2NH3  and  Cu(NH3)4S04. 

50.  Nitric  Acid  and  the  Oxides  of  Nitrogen. 

a.  Mix  10  grams  of  sodium  nitrate  (NaN03)  with  about  8  c.c. 
of  concentrated  sulfuric  acid  in  a  retort,  arrange  a.  receiving 
flask  which  can  be  cooled  by  running  water,  as  in  Exp.  35,  and 
warm  gently,  protecting  the  retort  with  a  wire  gauze,  until  the 
nitric  acid  has  been  entirely  distilled  over.     This  should  give 
about  5  c.c.  of  concentrated  nitric  acid.     Note  the  color  of 
the  acid  and  of  the  gas  above  it.     Pure  nitric  acid  is  colorless. 
Can  you  explain  the  color  here?     (If  this  acid  cannot  be  used 
at  once  dilute  it  with  one-half  its  volume  of  water  before  putting 
it  away.) 

b.  Place  about  1  c.c.  of  this  acid  in  a  test-tube,  and  dilute  it 
with  one-half  its  volume  of  distilled  water.     Moisten  a  piece 
of  filter  paper  with  this  solution,  then  allow  it  to  dry  out  on 
the  radiator.     Dilute  a  few  drops  of  the  acid  with  2-3  c.c.  of 
water  and  add  barium  chloride  (BaCl2)  solution.     Now  add  a 
small  amount  of  sulfur  to  a  new  portion  of  the  diluted  acid,  warm 
till  reaction  seems  to  cease,  noting  the  appearance  of  the  solution 
and  of  the  gas  evolved,  and  again  test  the  solution  with  filter 
paper  as  before  (see  Exp.  44).     Account  for  the  difference  in  the 
two  tests.     What  has  become  of  the  sulfur?     Add  a  few  drops 
of  barium  chloride  solution  to  the  mixture  in  the  test-tube 
(after  diluting  with  distilled  water)  and  explain  the  change. 
Does  nitric   acid  form  a  precipitate  with  barium  chloride? 
What  property  of  nitric  acid  does  this  reaction  with  sulfur 
illustrate? 

c.  Nitric  Oxide. — Arrange  a  generating  flask  as  usual,  with 
thistle  tube  and  delivery  tube.     Place  about  3  grams  of  copper 
in  the  flask  and  add  the  rest  of  the  nitric  acid  prepared  above, 
diluted  with  one-half  its  volume  of  water.     Heat  the  flask  and 
contents  and  collect  one  bottle  of  the  gas  evolved  by  displace- 


66  NITROGEN 

ment  of  water,  noting  the  color  of  the  gas  in  the  generating 
flask,  and  in  the  gas  collecting  bottle.  Account  for  the  differ- 
ence in  color.  At  the  same  time  collect  a  test-tube  one-third 
full  of  the  gas,  note  its  color,  and  while  the  mouth  of  the  tube  is 
still  under  water  pour  into  the  tube  (by  pouring  upward)  the 
air  from  another  test-tube  which  is  two-thirds  full  of  air,  the 
rest  of  the  tube  being  full  of  water  and  inverted  in  the  bath 
like  the  first  one.  Account  for  the  change  in  appearance  and 
volumes  of  the  gases  after  mixing.  Shake  the  tube  while  still 
inverted  under  water  to  determine  whether  the  gas  formed  is 
soluble  in  water.  Does  this  experiment  throw  any  light  on  the 
reason  for  the  difference  in  cclor  between  the  gas  in  the  generat- 
ing flask  and  in  the  collecting  bottle?  Write  the  reactions 
involved.  Will  the  nitric  oxide  support  the  combustion  of  a 
burning  splinter?  Of  phosphorus?  (Test  both  of  these  on  the 
bottle  of  gas.) 

d.  Nitrous  Oxide. — Fit  a  small  flask  with  stopper  and  delivery 
tube  leading  to  the  water  trough.  Place  about  10  grams  of 
ammonium  nitrate  (NH4N03)  in  the  flask  and  heat  gently 
(too  rapid  heating  may  cause  an  explosion)  collecting  two 
bottles  of  the  gas  evolved  by  displacement  of  water.  Note 
the  color  of  the  gas.  Does  it  support  the  combustion  of 
wood  and  phosphorus?  Compare  it  with  oxygen  in  this 
respect.  Is  there  any  apparent  reaction  with  the  oxygen  of 
the  air?  Compare  it  with  nitric  oxide  in  this  respect.  How 
could  a  given  sample  of  gas  be  tested  to  determine  whether  it 
was  nitric  oxide  or  nitrous  oxide? 

61.  Behavior  of  Nitrates  and  Nitrites  on  Heating. 

Heat  samples  of  sodium  nitrate,  sodium  nitrite  and  powdered 
lead  nitrate  in  hard  glass  test-tubes.  Note  the  color  of  the  gas 
evolved  and  test  it  for  oxygen.  Add  1  c.c.  of  dilute  hydrochloric 
acid  to  the  residue  after  heating,  in  each  case,  again  noting  the 
color  of  the  gas  evolved.  Explain  the  reactions  and  write 
equations.  Arrange  in  tabular  form  the  products  obtained  on 
heating  sodium  nitrate,  ammonium  nitrate,  ammonium  nitrite, 


OXIDATION  WITH  NITRIC  ACID  67 

sodium  nitrite  and  lead  nitrate,  and  the  products  obtained  when 
the  residue  from  heating  is  treated  with  dilute  hydrochloric 
acid. 

The  Combining  Weight  of  a  Metal  by  Oxidation  with  Nitric 
Acid. 

Combining  Weight  of  Tin.  Quantitative. — Carefully  clean  and 
dry  an  evaporating  dish  and  cover  glass,  weigh  it,  then  place 
in  it  about  1  gram  of  tin  foil  and  weigh  again.  Convert  the 
metal  to  its  oxide  by  adding  dilute  nitric  acid  (HN03),  about 
5  c.c.  at  a  time,  until  all  reaction  apparently  ceases.  (Note 
the  color  of  the  gas  evolved  and  explain.)  Then  remove  the 
cover  and  evaporate  on  a  steam  bath  to  remove  the  excess 
of  nitric  acid  and  water,  finally  heating  the  covered  evaporating 
dish  with  the  direct  flame  to  remove  the  last  traces  of  water. 
Weigh,  heat  again  and  reweigh  until  the  dish  and  contents 
have  come  to  constant  weight.  The  product  in  the  dish  is  an 
oxide  of  tin.  From  the  data  obtained  calculate  the  combining 
weight  of  the  tin.  If  any  nitrate  of  tin  was  formed  during 
the  process  what  has  become  of  it?  Was  there  any  indication 
that  a  nitrate  was  being  decomposed  while  the  dish  was  heated 
with  the  direct  flame?  Write  the  reactions  involved.  The 
formula  for  the  tin  oxide  may  be  determined  from  the  combin- 
ing proportions  and  the  atomic  weights  of  tin  and  oxygen. 


CHAPTER  IX 
THE  ATMOSPHERE 

53.  The  Determination  of  Oxygen  in  the  Air.     Quantitative. 

If  a  eudiometer  is  not  available  for  this  experiment  a  test-tube 
may  be  roughly  graduated  in  10  c.cv  marking  it  with  a  pencil 
for  writing  on  glass,  or  with  file  marks,  and  used  in  its  place. 


v 


FIG.  15. 

Invert  a  eudiometer  containing  the  air  of  the  laboratory  over  a 
\vater  bath,  clamp  it  in  position,  read  the  volume  of  air  at  at- 
mospheric pressure,  and  insert  into  the  air  a  bit  of  yellow 
phosphorus  fastened  on  the  end  of  a  wire.  Fig.  15.  The 
phosphorus  may  be  prepared  by  melting  a  piece  of  yellow 
phosphorus  about  the  size  of  a  pea,  under  water,  placing  the 
wire  in  the  melted  phosphorus  and  allowing  the  phosphorus  to 
solidify  in  this  position.  Yellow  phosphorus  should  be  kept  under 
water  and  should  not  be  handled  with  the  bare  hands,  as  it  causes 
very  unpleasant  burns.  Allow  this  to  stand  until  there  is  no 
further  diminution  in  volume  of  the  gas  and  again  read  the  volume 

68 


COMPOSITION  OF  AIR  69 

at  atmospheric  pressure.  Will  the  gas  remaining  in  the  tube 
support  combustion?  What  is  it?  What  is  the  per  cent  by 
volume  of  oxygen  in  the  air?  In  this  experiment  the  air  was 
always  saturated  with  moisture,  and  was  measured  at  at- 
mospheric temperature  and  pressure.  Is  it  necessary  to  calcu- 
late the  volume  of  the  dry  gas  under  standard  conditions,  before 
and  after  absorption  of  the  oxygen,  before  calculating  its  per 
cent  by  volume?  Would  it  be  necessary  to  do  so  in  order  to 
calculate  its  per  cent  by  weight? 

54.  The  Composition  of  Air  Dissolved  in  Water.    Quantitative. 

The  composition  of  air  dissolved  in  water  may  be  very  easily 
determined  by  the  method  used  for  the  determination  of  oxygen 
in  atmospheric  air,  the  only  change  in  manipulation  being  the 
source  of  the  sample  for  analysis.  Use  the  apparatus  shown  in 
Fig.  16,  in  which  the  eudiometer  or  graduated  test-tube  is  fitted 
with  a  two  hole  stopper,  one  hole  of  which  carries  a  tube  leading 
just  to  the  lower  surface  of  a  cork  fitted  to  a  500  c.c.  Erlenmeyer 
flask,  while  the  other  hole  is  fitted  with  a  tube  bent  so  that  it 
leads  under  the  surface  of  water  in  a  beaker.  Each  of  these  con- 
necting tubes  must  be  in  two  pieces,  connected  with  a  rubber  tube 
which  can  be  closed  with  a  pinch  cock.  Fill  the  whole  apparatus 
completely  with  fresh  tap  water  and  clamp  the  eudiometer  in 
an  upright,  inverted  position  above  the  Erlenmeyer  flask,  open 
both  the  pinch  cocks  (if  water  runs  out  on  opening  the  cocks  it 
is  an  indication  of  a  leak,  which  must  be  corrected  before 
proceeding)  and  slowly  heat  the  water  in  the  Erlenmeyer  flask. 
The  gas  which  separates  will  rise  to  the  top  of  the  eudiometer, 
while  a  small  amount  of  water  will  run  out  into  the  beaker  be- 
cause of  the  expansion  of  the  water  on  heating  and  its  dis- 
placement by  the  gas.  The  water  may  be  brought  to  the  boiling- 
point,  but  it  must  not  be  allowedlo  boil  lest  steam  escape,  carry- 
ing with  it  some  of  the  gases.  If  this  does  not  supply  a  sufficient 
sample  of  gas  for  analysis  the  two  pinch  cocks  may  be  closed, 
the  Erlenmeyer  flask  refilled  with  fresh  tap  water,  taking  care 
that  it  is  completely  filled  with  water  when  the  eudiometer  is 


70 


THE  ATMOSPHERE 


replaced,  and  the  procedure  repeated  until  a  sufficient  sample  is 
obtained.  When  a  sufficient  sample  of  the  gas  is  obtained  (20 
c.c.  will  be  enough  for  rough  results)  close  both  the  pinch  cocks, 
and  bring  the  eudiometer,  with  its  stopper  and  connecting  tubes, 
under  the  surface  of  water  in  a  water  bath,  then  remove  the 
stopper  and  read  the  volume  after  the  gas  has  come  to  at- 
mospheric temperature,  and  proceed  to  determine  the  oxygen  as 


FIG.  16. 

in  Exp.  53.  Calculate  the  per  cent  composition  by  volume  of 
the  gas  analyzed,  and  compare  with  the  per  cent  composition 
calculated  from  the  solubility  of  the  two  gases  and  their  partial 
pressures  in  the  air.  What  would  be  the  relative  amounts 
dissolved  if  the  air  were  40  per  cent  oxygen?  State  Henry's 
law. 


CHAPTER    X 
PHOSPHORUS,  ARSENIC,  ANTIMONY  AND  BISMUTH 

55.  The  Oxides  of  Phosphorus. 

Connect  a  hard  glass  combustion  tube  open  at  each  end  with  a 
delivery  tube  fitted  with  a  pinch  cock  and  leading  to  the  bottom 
of  a  test-tube  or  small  flask  one-third  full  of  water  which  is  fitted 
with  a  two  hole  stopper,  the  other  hole  of  which  carries  a  tube 
connected  with  a  suction  pump  or  aspirator  bottle  (Fig.  17). 


FIG.  17. 

a.  Place  red  phosphorus  the  size  of  a  pea  in  the  hard  glass 
tube,  light  by  touching  with  a  red  hot  file,  and  by  manipulating 
the  pinch  cock  keep  the  phosphorus  burning  quietly  while  air  is 
drawn  over  it  and  the  products  of  combustion  are  absorbed^by 
the  water  in  the  test-tube.  Rinse  out  the  combustion  tube  and 
the  delivery  tube  into  the  solution  in  the  test-tube,  filter  and  test 
the  filtrate  with  litmus.  Make  the  following  tests  on  this 
solution : 

71 


72       PHOSPHORUS,  ARSENIC,  ANTIMONY,  ETC. 

b.  To  a  small  portion  add  a  few  drops  of  potassium  perman- 
ganate solution  and  warm. 

c.  To  another  portion  add  silver  nitrate  solution,  then  sodium 
bicarbonate  (NaHC03)  till  a  precipitate  becomes  permanent, 
then  warm. 

d.  Add  a  small  portion  to  albumin  solution  acidified  with 
acetic  acid.     What  do  each  of  these  tests  indicate?     What 
materials  are  thus  shown  to  be  present  in  the  solution? 

e.  Evaporate  the  rest  of  this  solution  in  a  porcelain  dish,  adding 
nitric  acid  till  brown  fumes  cease  to  be  given  off.  (What  are  they  ? 
Explain  their  formation.)  Evaporate  till  acid  fumes  are  no 
longer  given  off,  then  heat  with  the  direct  flame  until  there  is 
no  further  change,  cool,  and  dissolve  in  water.  Test  this  solu- 
tion with  potassium  permanganate  and  silver  nitrate  solutions 
as  above,  and  explain  the  results.  What  is  now  present  in  the 
solution?  Allow  the  solution  to  stand  for  several  days  and  test 
again  with  silver  nitrate.  Explain. 

56-  The  Preparation  of  Primary,  Secondary  and  Tertiary 
Orthophosphates,  and  their  Decomposition  Products  on  Heat- 
ing. Quantitative. 

a.  Arrange  twelve  test-tubes  in  groups  of  three,  and  to  the 
first  of  each  group  add  a  solution  of  primary  sodium  phosphate 
(NaH2P04);    to    the    second,    secondary    sodium    phosphate 
(Na2HP04);  to  the  third,  tertiary  sodium  phosphate  (Na3P04), 
and  label  each  tube.     To  each  of  the  tubes  in  the  first  group  add 
a  few  drops  of  alizarine  green,  to  the  second  group  add  phenol- 
phthalein,  to  the  third  congo  red,  and  to  the  fourth  methyl  orange. 
Record  the  colors  formed  in  tabular  form,  noting  at  the  same 
time  the  colors  of  each  indicator  in  a  solution  of  a  strong  acid 
and  of  a  strong  alkali.     Which  indicator  should  be  used  in  the 
preparation  of  each  of  the  sodium  salts  of  orthophosphoric 
acid  by  neutralization  of  the  acid  with  sodium  hydroxide? 
Get    the    approval    of   an   instructor    on   this   point    before 
proceeding. 

b.  Prepare  100  c.c.of  an  approximately  formular  solution  of 
phosphoric  acid  by  diluting  the  concentrated  acid  (see  page  7 


ARSENIC  73 

for  the  table  of  per  cent  composition  and  density)  and  an  equal 
amount  of  formular  sodium  hydroxide  solution  from  the  solid 
sticks. 

c.  Into  each  of  three  evaporating  dishes  measure  accurately 
(pipette)  10  c.c.  of  the  acid  solution  after  it  has  been  well  mixed, 
then  neutralize  the  acid  in  each  dish  with  the  sodium  hydroxide 
solution  measured  from  a  burette,  and  using  the  selected  in- 
dicators, so  that  the  primary,  secondary  and  tertiary  salts  will 
be  formed.     Record  the  volumes  of  alkali  required  in  each  case. 
What  ratio  do   these  amounts  bear   to   each   other?     What 
ratio  should  they  bear,    theoretically?     Evaporate  the  solu- 
tions   on  the  water  bath,  then  heat  the  dry  salts  with  the 
direct  flame  until   all   moisture   is   removed.     What   salt   is 
now   present   in   each   case?     Dissolve   each   of   these   three 
salts  in  water  and  test  each  solution  with  silver  nitrate  solution, 
and  with  albumin  solution  acidified  with  acetic  acid.     Do  these 
tests  confirm  your  conclusions  regarding  the    salts  present? 

d.  Devise  and  perform  tests  to  show  that  most  phosphates  are 
insoluble  in  water  but  soluble  in  acids. 

57.  Arsenic. 

a.  Place  about  a  gram  of  arsenopyrites  in  the  center  of  a 
hard  glass  tube  such  as  was  used  in  preparing  the  oxides  of 
phosphorus  (Exp.  55),  having  one  end  closed  with  a  cork,  and 
the  other  carrying  a  small  bore  delivery  tube.  Heat  the  pyrite, 
when  a  sublimate  will  collect  on  either  side  of  the  hot  part  of 
the  tube.  What  is  it?  After  a  considerable  sublimate  has 
collected  allow  the  tube  to  cool,  remove  the  residue  of  pyrite 
and  connect  the  combustion  tube  with  a  test-tube  or  flask,  and 
suction  as  in  Exp.  55,  Fig.  17,  and  burn  one  of  the  two  deposits 
(save  the  other  for  c)  collecting  the  products  in  the  test-tube 
as  before.  Compare  the  solubility  of  arsenic  oxides  with  that 
of  phosphorus  oxides.  Filter  the  solution  of  the  oxides  (keep 
the  insoluble  residue  for  b)  and  test  a  small  portion  of  the  fil- 
trate for  sulfur  dioxide  by  adding  hydrochloric  acid  then  barium 
chloride — a  precipitate  here  would  indicate  what  substance  ? 


74       PHOSPHORUS,  ARSENIC,  ANTIMONY,  ETC. 

If  any  precipitate  appears  filter,  then  add  potassium  perman- 
ganate solution  to  the  filtrate  and  heat.  A  precipitate  here  is 
due  to  what  substance?  Explain  its  presence. 

b.  Wash  the  residue  of  the  arsenic  oxide  several  times  with 
distilled  water,  then  punch  a  hole  in  the  bottom  of  the  filter 
paper  and  wash  the  oxide  into  a  beaker  with  20-30  c.c.  of  dis- 
tilled water.  Heat  this  to  boiling  to  dissolve  as  much  of  the 
oxide  as  possible,  filter  if  the  solution  is  not  clear  and  test  the 
solution  for  sulfur  dioxide  as  before.  Is  a  precipitate  of  barium 
sulfate  formed?  Is  the  potassium  permanganate  decolorized? 
Explain.  To  a  small  portion  of  the  solution  add  silver  nitrate 
then  sodium  carbonate  solution  till  a  permanent  precipitate 
forms  (one  drop  is  usually  enough)  and  boil.  Is  the  precipi- 
tate silver  carbonate?  Devise  and  perform  an  experiment  to 
answer  this  question.  What  is  the  precipitate  and  what  is  its 
reaction  on  boiling?  Compare  with  the  similar  compound  of 
phosphorus. 

Into  the  remainder  of  the  solution  pass  hydrogen  sulfide  from 
a  clean  delivery  tube  (if  the  containers  and  solution  have  not 
been  kept  perfectly  clean  and  free  from  electrolytes  this  will 
be  a  failure)  and  note  the  change  in  color.  This  should  give 
a  colloidal  solution  of  arsenic  sulfide.  Compare  the  physical 
properties  of  this  solution  with  those  of  a  true  solution.  Divide 
the  solution  into  two  portions — to  one  add  hydrochloric  acid, 
to  the  other  sodium  chloride  solution  and  explain  the  change. 
Would  it  be  classified  as  a  chemical  or  physical  change? 

c.  Dissolve  the  other  portion  of  the  arsenic  sublimate  (a)  in 
concentrated  nitric  acid  and  evaporate  till  the  excess  of  acid  is 
removed  (hood),  add  water  and  to  a  small  portion  of  the  solu- 
tion, filtered  if  not  clear,  add  silver  nitrate  and  sodium  car- 
bonate solutions  till  a  permanent  precipitate  forms.  Why  is 
the  sodium  carbonate  necessary?  What  is  the  precipitate? 
Boil.  Compare  with  the  silver  arsenite  precipitate  prepared  in 
6.  Evaporate  the  rest  of  the  solution  to  dryness  and  heat  till 
all  moisture  is  removed.  What  remains?  Compare  with  the 
effect  of  heating  phosphoric  acid. 


BISMUTH  75 

58.  Antimony. 

Treat  1-2  grams  of  stibnite  (what  is  it  ?)  with  concentrated 
hydrochloric  acid,  warming  in  a  covered  evaporating  dish  until 
the  reaction  becomes  slow.  Note  the  odor  of  the  gas  evolved. 
What  is  it?  Remove  the  solution  from  the  excess  of  stibnite 
and  add  an  equal  volume  of  water.  If  a  white  precipitate 
forms  (what  is  it?)  add  hydrochloric  acid  till  it  dissolves.  If 
the  solution  turns  yellow  or  a  yellow  precipitate  appears  (due 
to  what?)  boil  till  the  color  disappears.  Make  this  solution 
alkaline  with  ammonium  hydroxide,  filter  out  the  precipitate 
and  wash  with  distilled  water.  What  is  the  precipitate?  Boil 
a  small  portion  of  the  precipitate  with  distilled  water  and  test 
with  litmus.  Is  the  hydroxide  soluble?  To  two  other  por- 
tions of  the  precipitate  add  sodium  hydroxide  solution  and 
hydrochloric  acid  respectively.  Draw  conclusions  as  to  the 
acidic  or  basic  character  of  the  hydroxide.  Compare  with  the 
character  of  the  oxides  of  arsenic  and  phosphorus. 

59.  Bismuth. 

Add  2-3  c.c.  concentrated  nitric  acid  to  about  0.5  gram  of 
metallic  bismuth,  boil  the  solution  till  nitric  oxide  fumes  are 
no  longer  evolved  and  allow  to  crystallize  by  cooling  and  spon- 
taneous evaporation.  Heat  a  small  portion  of  the  crystals  in 
a  dry  test-tube,  noting  the  gas  evolved.  Add  a  second  small 
portion  to  water  and  test  with  litmus.  Account  for  the  acid 
solution  and  for  the  insoluble  substance.  Add  nitric  acid  to 
this  until  solution  is  complete,  then  prepare  bismuth  hydroxide 
by  adding  ammonium  hydroxide  to  the  solution.  Filter  and 
wash  once  with  distilled  water.  Treat  part  of  the  precipitate 
with  a  25  per  cent  solution  of  sodium  hydroxide,  and  part  with 
dilute  hydrochloric  acid.  Is  bismuth  hydroxide  basic  or  acidic? 
Indicate  its  ionization  products  and  the  effect  of  the  addition 
of  sodium  hydroxide  or  hydrochloric  acid  on  the  ionization. 


CHAPTER  XI 

CARBON 

60.  Carbon. 

a.  Fill  a  crucible  about  one-fourth  full  of  cane  sugar,  and  heat 
until  the  mass  solidifies  while  hot.  testing  from  time  to  time  to 
determine  whether  the  gas  evolved  burns.     Heat  4-5  grams  of 
copper  oxide  in  a  crucible  to  remove  water,  and  when  cool  mix 
with  some  of  the  sugar  charcoal  in  a  dry,  hard  glass  test-tube 
and  heat,  while  watching  for  the  condensation  of  moisture  in 
the  cool  part  of  the  tube.     Pass  some  of  the  gases  evolved  into 
lime  water,  and  explain  the  result.     To  what  element  in  the 
charcoal  is  the  moisture  due?     Is  charcoal  often  pure  carbon? 
Name  the  forms  in  which  pure  carbon  occurs.     Burn  the  rest 
of  the  charcoal,  observing  the  rate  and  ease  of  combustion. 
Is  carbon  an  active  element?     Does  it  burn  with  a  flame? 

b.  Heat  gently  about  2  grams  of  soft  coai  in  a  crucible  covered 
with  an  inverted  funnel,  noting  the  formation  of  a  gas  (does  it 
burn?)    and  of  tar.     When  combustible  gases  are  no  longer 
given  off  cool  and  observe  the  appearance  of  the  coke  formed. 
Is  it  pure  carbon?     To  answer  this  apply  the  test  used  in  a 
on  one  portion  of  the  coke  and  place  the  crucible  containing  the 
rest  of  the  coke  in  an  inclined  position  and  heat  until  the  carbon 
has  all  burned  away.     What  is  left?     What  are  some  of  the 
points  to  be  considered  in  buying  coal  for  heating  purposes? 

61.  Carbon  Dioxide. 

Add  dilute  hydrochloric  acid  to  marble  in  a  generating  flask 
(Fig.  7)  and  pass  the  gases  evolved  into  lime  water.  What  is 
formed?  Attach  a  wash  bottle  to  the  generator  so  that  the  gas 
bubbles  through  water  (to  remove  hydrochloric  acid  which 
might  be  carried  over  mechanically),  then  pass  the  gas  into  a 

76 


CARBON  MONOXIDE  77 

flask  containing  about  100  c.c.  of  water  until  the  solution  is  satu- 
rated. Test  the  solution  with  litmus  paper,  taste  it.  What  is 
present  ?  Keep  this  solution  (stoppered)  for  Exp.  64.  Determine 
whether  the  gas  is  lighter  or  heavier  than  air,  and  fill  a  bottle  of 
the  gas  by  air  displacement.  Insert  into  this  a  burning  taper. 
Pour  the  gas  from  one  bottle  into  another  and  demonstrate  that 
it  has  really  passed  from  one  to  the  other.  Direct  the  stream 
of  gas  from  the  generator  upon  the  flame  from  a  burning  match 
or  taper,  and  explain.  What  commercial  application  is  made 
of  this  property? 

62.  Determination  of  the  Density  of  Carbon  Dioxide.     Quan- 
titative. 

Which  of  the  methods  previously  used  for  density  determina- 
tions can  be  applied  to  carbon  dioxide?  Arrange  the  apparatus 
and  proceed  with  the  determination  after  getting  the  approval 
of  an  instructor. 

63.  Carbon  Monoxide.     Poison.      Do  not   Let   the   Gas  Escape 
into  the  Room. 

a.  Mix  about  3  grams  of  solid  oxalic  acid  with  an  equal 
volume  of  concentrated  sulfuric  acid  in  a  generating  flask  and 
heat,  collecting  three  bottles  of  the  gas  by  displacement  of  water. 
To  stop  the  evolution  of  gas  allow  the  generating  flask  to  cool. 

b.  Add  lime  water  to  the  first  bottle  and  shake.     What  is 
formed? 

c.  Invert  the  second  bottle  as  quickly  as  possible  over  a  beaker 
of  dilute  potassium  hydroxide  solution  and  allow  to  stand  until 
absorption  is  complete,   then  cover,   set  the  bottle  upright 
without  pouring  out  the  water  that  was  absorbed  and  apply  a 
light.     What  is  this  gas?     What  were  the  relative  proportions 
of  carbon  dioxide  and  carbon  monoxide  in  the  gas  mixture? 

d.  Apply  a  light  to  the  third  bottle  of  the  gas  mixture,  and 
explain  the  result. 

64.  Hard  Water. 

a.  Dilute  some  lime  water  with  an  equal  volume  of  distilled 
water  and  pass  in  carbon  dioxide  until  the  precipitate  first 


78  CARBON 

formed  (what  is  it?)  redissolves  (what  is  formed?).  Filter 
the  solution  if  it  is  not  entirely  clear,  and  divide  into  five  por- 
tions. Boil  one  portion.  To  the  other  four  add  sodium  car- 
bonate, lime  water,  soap  and  alum  solutions  respectively.  Draw 
conclusions  as  to  the  efficiency  of  these  reagents  and  of  boiling 
as  softening  agents? 

b.  Pour  20  c.c.  of  the  carbonic  acid  solution  prepared  in 
Exp.  61  over  a  small  amount  of  powdered  marble  or  chalk,  and 
allow  to  stand  over  night,  then  filter  and  test  the  filtrate  by 
boiling, with  sodium  carbonate,  etc.  What  was  present  in  the 
solution?  What  is  the  importance  of  this  reaction  in  nature? 

65.  Soap. 

a.  Dissolve  about  0.5  gram  of  soap  by  heating  in  25  c.c.  of 
distilled  water  and  test  the  solution  with  litmus  and  explain. 
Add  solutions  of  sodium  and  calcium  salts  to  different  portions 
of  this  soap  solution.     Which  would  be  more  objectionable  in 
water  for  laundry  purposes,  sodium  or  calcium  bicarbonate? 

b.  Place  equal  quantities  of  animal  charcoal  in  each  of  two 
test-tubes;  to  one  add  10  c.c.  of  hot  soap  solution,  to  the  other 
10  c.c.  of  hot  water,  filter  each  and  compare  the  filtrates.     The 
soap  solution  tends  to  combine  with  the  carbon  to  form  an 
adsorption  compound  which  reduces  the  tendency  for  the  carbon 
to  stick  to  other  materials,  here  the  filter  paper.    This  property 
partly  explains  the  action  of  soap  in  cleansing. 

c.  To  10  c.c.  of  soap  solution  add  dilute  hydrochloric  acid  to 
acid  reaction.     Warm  gently  until  the  white  precipitate  melts 
and  floats  on  the  surface.     What  is  it? 

66.  Methane. 

a.  Grind  together  one  part  of  sodium  acetate  with  four  parts 
of  soda  lime  (quick  lime  and  sodium  hydroxide)  and  place  the 
mixture  in  a  hard  glass  test-tube  fitted  with  a  delivery  tube. 
Incline  the  tube  so  that  water  driven  from  the  contents  on  heat- 
ing will  not  run  back  upon  the  hot  part  of  the  glass,  and  warm, 
collecting  the  gas  evolved  by  displacement  of  water. 


HYDROCARBONS  79 

b.  Shake  one   bottle   of   the   gas  with  a  few  cubic    centi- 
meters of  potassium  permanganate  solution. 

c.  Shake  another  with  bromine  water. 

d.  Burn  a  third  and  test  the  products  of  combustion  remain- 
ing in  the  bottle  for  carbon  dioxide. 

e.  Light  a  few  bubbles  of  the  gas  as  they  escape  from  the 
water  and  note  the  nature  of  the  flame. 

67.  Ethylene. 

Add  9  c.c.  alcohol  slowly  to  60  c.c.  concentrated  sulfuric  acid, 
stirring  constantly.  Pour  this  mixture  into  a  flask  fitted  with 
a  delivery  tube  leading  under  water,  and  clamped  so  as  to  rest 
on  a  sand  bath.  Add  a  little  sand  to  reduce  foaming  and  heat 
gentry  until  there  is  a  steady  evolution  of  gas  (avoid  over- 
heating as  the  mixture  foams  considerably)  and  collect  bottles 
of  the  gas,  repeating  all  the  tests  applied  to  methane  and 
comparing  the  results. 

68.  Acetylene. 

a.  Place  15  grams  of  calcium  carbide  in  a  flask  fitted  with  a 
separatory  funnel  and  delivery  tube  which  is  connected  with  a 
hard  glass  combustion  tube  open  at  each  end  and  fitted  with  a 
delivery  tube  leading  under  water,  Fig.  18.  Generate  the  gas 
by  allowing  water  from  the  separatory  funnel  to  drop  upon  the 
calcium  carbide,  and  collect  bottles  of  the  gas  to  which  all  the 
tests  under  methane  should  be  applied.  Compare  with  methane 
and  ethylene. 

6.  After  all  the  air  has  been  driven  out  from  the  generating 
flask  and  delivery  tubes  heat  the  hard  glass  tube  while  still 
generating  acetylene,  and  observe  the  deposit  on  the  glass,  and 
the  changed  appearance  of  the  gas  evolved,  and  explain. 

69.  The  Oxidation  Products  of  Alcohol. 

Add  2  c.c.  of  concentrated  sulfuric  acid  to  a  mixture  of  8  c.c. 
water  and  2  c.c.  50  per  cent  alcohol,  mix  thoroughly  then  add 
a  small  crystal  of  potassium  dichromate  and  warm.  Test  the 


80 


CARBON 


gas  evolved  by  holding  in  it  strips  of  moist  litmus  paper;  and  note 
its  odor.     This  is  acetaldehyde. 

To  1  to  2  c.c.  of  concentrated  sulfuric  acid  add  about  5  drops  of 
50  per  cent  alcohol,  mix,  then  add  a  small  crystal  of  potassium  di- 


ll 


FIG.  18. 


chromate  and  warm  till  the  reaction  begins  (care !) .  Test  the  gas 
evolved  with  moist  litmus  paper  and  note  its  odor.  What  is  it? 
Write  equations  for  both  reactions  and  explain  the  change  in 
color  of  the  solutions. 


CHAPTER  XII 

SILICON 

70.  Silicon  Tetrafluoride. 

Grind  together  1  gram  of  silicon  dioxide  (sand)  and  3  grams 
of  calcium  fluoride,  moisten  with  concentrated  sulfuric  acid  and 
heat  in  a  test-tube  fitted  with  a  delivery  tube  leading  just  above 
the  surface  of  water.  What  is  the  precipitate  formed  on  the 
surface  of  the  water?  Explain  the  fumes  above  the  water. 
Test  the  water  solution  by  filtering,  then  adding  to  the  filtrate 
dilute  hydrochloric  acid  and  barium  chloride  solution,  and  warm- 
ing. What  is  formed?  Write  all  reactions. 

71.  Silicic  Acid. 

a.  Grind  together  in  an  iron  crucible  1  gram  of  sand  and  3 
grams  of  sodium  hydroxide  (solid)  and  heat  until  the  mass 
which  at  first  fuses,  turns  solid.     Cool,  dissolve  in  25  c.c.  water 
and  filter.     To  one  portion  of  the  filtrate  (what  is  it?)  add  con- 
centrated or  dilute  hydrochloric  acid  very  slowly.     Pour  another 
portion,  about  6  c.c.,  into  a  test-tube  containing  about  6  c.c.  of 
concentrated  hydrochloric  acid.     Has  there  been  any  difference 
in  the  chemical  reaction  in  these  two  tubes?     Explain    the 
difference  in  appearance.     Evaporate  part  of  the  clear  solution 
of  silicic  acid  in  a  porcelain  dish  and  heat  the  dry  residue  with 
the  direct  flame,  then  cool  and  determine  whether  it  is  still  com- 
pletely soluble  in  water. 

b.  Dialyze  the  rest  of  the  silicic  acid  solution  as  follows: 
Prepare  a  collodion  sack  for  dialysis  by  wetting  the  inside  walls 
of  a  50  c.c.  beaker  which  has  been  carefully  cleaned  and  dried, 
with  a  solution  of  collodion  in  ether  and  allowing  the  ether  to 
evaporate,  then  when  the  odor  of  ether  has  disappeared  remove 
the  sack  from  the  beaker  (loosen  the  edge  of  the  collodion  from 

81 


82 


SILICON 


the  glass  with  a  knife,  then  the  sack  may  be  withdrawn  with  the 
finger,  or  by  allowing  water  to  run  in  between  the  glass  and  the 
collodion)  and  fill  with  the  solution  of  colloidal  silicic  acid. 
Suspend  this  sack  in  a  beaker  of  distilled  water  so  that  the  level 
of  the  water  is  slightly  lower  than  the  level  of  the  solution  in  the 
sack  (Fig.  19)  and  change  the  dialysate  (the  solution  in  the 


FIG.  19. 

beaker,  consisting  of  materials  which  have  dialyzed  through 
the  membrane)  frequently  until  it  gives  only  a  very  faint  test 
for  chlorides  after  standing  in  contact  with  the  sack  for  about 
an  hour.  Test  portions  of  the  dialyzed  solution  (the  solution 
remaining  in  the  sack)  with  barium  chloride  and  copper  sulfate 
solutions.  Would  barium  or  copper  silicates  be  soluble  if 
formed  ?  Is  the  silicic  acid  ionised  in  this  solution  ?  Evaporate 
another  portion  to  dryness,  heat  with  the  direct  flame  and 
determine  whether  the  residue  is  soluble  in  water.  Has  it 
undergone  a  chemical  change? 


CHAPTER  XIII 
BORON 

72.  Preparation  of  Boric  Acid  from  Borax. 

Saturate  50  c.c.  of  distilled  water  with  borax,  test  the  solution 
with  litmus  and  account  for  the  result.  Add  dilute  sulfuric  acid 
to  this  solution  until  no  further  precipitate  forms,  then  cool  and 
filter  the  cold  solution.  Wash  once  or  twice  with  cold  distilled 
water  then  purify  the  precipitate  by  redissolving  in  the  least 
possible  amount  of  boiling  water  and  cool  under  running  water. 
Filter  and  wash  the  precipitate  with  cold  water  till  the  wash 
water  is  free  from  sulfates,  then  dry  thoroughly  by  pressing 
between  folds  of  filter  paper  and  then  warming  on  the  radiator 
or  in  a  drying  oven  at  a  temperature  below  100°  C.  What  is  the 
residue?  Dissolve  some  in  water  and  test  the  solution  with 
litmus.  Mix  a  small  portion  with  an  equal  quantity  of  sodium 
chloride  and  heat  the  mixture  in  a  test-tube.  What  is  formed? 
Heat  a  small  portion  in  a  dry  test-tube.  Do  these  tests  prove 
whether  the  material  is  an  acid  or  an  acid  anhydride? 

73.  Boric  Acid  Anhydride   or  Boron  Trioxide  from  the  Acid. 
Quantitative. 

Carefully  clean  and  weigh  a  porcelain  crucible  and  place  in  it 
about  1  gram  of  the  well  dried  acid  from  the  previous  experi- 
ment (Exp.  72)  and  weigh  accurately.  Cover  the  crucible  and 
heat,  at  first  gently  and  then  with  the  full  heat  of  a  bunsen  burner. 
Cool  and  weigh,  and  repeat  until  the  crucible  and  contents  come 
to  constant  weight.  Calculate  the  per  cent  of  water  in  the  acid, 
and  the  number  of  molecules  of  water  to  one  of  the  acid 
anhydride. 

74.  The  Preparation  of  Metallic  Boron. 

Heat  some  of  the  prepared  boric  acid  (Exp.  72)  in  a  porcelain 
crucible  with  the  direct  flame  till  all  water  is  removed,  then  mix 

83 


84  BORON 

about  1  gram  of  the  anhydride  with  1.5  grams  of  magnesium 
powder.  Place  a  small  portion  of  this  mixture  in  a  crucible 
and  heat  till  glowing  ceases,  then  add  gradually  the  rest  of  the 
mixture,  heating  continuously.  Grind  the  product  in  a  mortar, 
wash  it  with  dilute  hydrochloric  acid  then  with  hot  concentrated 
hydrochloric  acid  diluted  1 :1  with  water,  and  finally  with  dis- 
tilled water  until  the  filtrate  comes  through  brown,  due  to  the 
formation  of  a  colloidal  solution  of  boron,  and  indicating  that 
electrolytes  are  not  present  in  appreciable  amounts.  Test  the 
solubility  of  this  boron  in  concentrated  nitric,  hydrochloric  and 
sulfuric  acids.  Fuse  a  small  portion  with  an  equal  quantity  of 
solid  potassium  hydroxide.  What  is  formed?  Dissolve  the 
fused  mass  in  as  little  water  as  possible  and  acidify  with  dilute 
sulfuric  acid.  What  is  formed?  How  may  tertiary  sodium 
borate  (Na3B03)  be  prepared?  Is  it  stable  in  solution? 

75.  Berates  of  the  Heavy  Metals. 

Will  borax  or  boric  acid  precipitate  copper  borate  from  copper 
sulfate  solution?  Is  copper  borate  soluble  in  water?  To 
answer  this  question  prepare  a  borax  bead  by  fusing  borax  on 
the  loop  of  a  platinum  wire,  dip  the  bead  into  copper  oxide  and 
heat  until  a  transparent  colored  bead  is  formed.  If  the  bead 
does  not  become  clear  after  heating  several  moments  it  is  be- 
cause too  much  copper  oxide  has  been  added.  What  copper 
compound  is  present  in  the  bead?  Break  up  the  bead  and  test 
its  solubility  in  water.  Boil  the  water  supension  and  note  the 
change  in  color  of  the  precipitate,  which  is  due  to  the  formation 
of  copper  oxide  from  the  hydroxide  first  formed.  Explain 
the  presence  of  copper  hydroxide  on  treating  the  bead  with 
water,  on  the  basis  of  the  ionization  theory. 


CHAPTER  XIV 

METALS  AND  NON-METALS 

76.  Hydrolysis. 

Dissolve  crystals  of  sodium  chloride,  potassium  chloride, 
ammonium  chloride,  cupric  chloride,  ferric  chloride,  stannous 
chloride  and  phosphorus  pentachloride  each  in  a  separate 
beaker  in  about  5  to  10  c.c.  of  water.  Test  the  solutions  with 
litmus,  then  boil  each,  covering  the  beaker  with  a  watch  glass 
on  which  is  placed  moist  red  and  blue  litmus  paper  so  that  it 
will  come  in  contact  with  the  vapor  from  the  boiling  solution. 
Note  and  account  for  the  effect  of  the  vapor  on  the  litmus  in 
each  case,  then  allow  the  solutions  to  evaporate  to  dryness  on 
the  water  bath,  dissolve  the  dry  residue  in  very  dilute  nitric 
acid  and  test  the  solution  for  chlorides.  Which  of  these  salts 
have  undergone  no  hydrolysis?  Which  have  been  partially 
hy droly zed ?  Which  have  been  completely  hy droly zed  ?  Define 
hydrolysis.  What  is  the  effect  of  water  on  the  chlorides  of 
metals  ?  Of  non-metals  ?  Do  the  results  of  this  experiment  lead 
to  the  conclusion  that  there  is  a  sharp  distinction  between 
metals  and  non-metals? 

77.  lonization. 

a.  Make  up  100  c.c.  each  of  thrice  normal  solutions  of  sulfuric, 
hydrochloric,  acetic,  phosphoric  and  oxalic  acids.  To  deter- 
mine the  relative  strength,  or  degree  of  ionization  of  these  acids 
place  5  c.c.  portions  of  each  acid  in  separate  test-tubes, 
label  each  tube  and  place  them  all  in  a  beaker  containing 
water,  so  that  they  will  not  suffer  any  great  change  of  tempera- 
ture. Now  into  each  tube  drop  a  small  granule  of  zinc  and  note 
the  rate  of  the  reaction,  j  udging  by  the  rate  of  evolution  of  hy- 
drogen in  each  tube.  Name  the  acids  in  the  order  of  their 

85 


86  METALS  AND  NON-METALS 

strength,  then  heat  the  water  bath  till  it  reaches  about  50°  and 
again  note  the  rate  of  the  reaction,  and  record.  If  the  weakest 
reaction  were  considered  the  standard  or  unit  of  strength, 
what  values  would  you  give  the  other  acids?  Compare  your 
values  with  those  in  the  table  of  degrees  of  ionization.  Repeat 
this  experiment  using  small  chunks  of  marble  instead  of  zinc, 
and  arrange  the  acids  according  to  their  strength,  as  before. 
How  does  this  order  agree  with  that  arranged  with  reference  to 
their  action  on  zinc? 

b.  Make  a  saturated  solution  of  potassium  chromate.  and 
place  equal  quantities  in  each  of  six  test-tubes.     Now  add  to  one 
of  the  test-tubes,  drop  by  drop,  the  acid  found  to  be  weakest  in 
the  previous  test,  until  a  just  noticeable  change  in  color  has 
occurred,  then  add  to  each  of  the  other  test-tubes  an  equal  num- 
ber of  drops  of  each  of  the  acids,  label  each,  and  to  the  sixth 
tube  add  an  equal  number  of  drops  of  concentrated  sulfuric  acid 
(this  represents  the  greatest  possible  change) .     Shake  each  tube 
and  arrange  in  the  order  of  their  depth  of  color,  and  compare 
with  the  order  of  the  acids  with  respect  to  their  activity  on  zinc 
or  marble. 

c.  Repeat  this  procedure,  using  a  solution  of  6  grams  po- 
tassium iodide  and  1  gram  potassium  bromate  per  liter  instead 
of  the  solution  of  potassium  chromate.     In  this  case  each  of  the 
tests  will  finally  come  to  the  same  color,  but  the  weakest  acid 
will  bring  about  the  change  most  slowly,  so  that  they  may  be 
arranged  according  to  either  of  two  variables,  i.e.,  the  depth  of 
color  at  a  fixed  time  after  adding  the  acid,  or  the  length  of 
time  required  for  each  tube  to  come  to  a  definite  depth  of  color. 
Record  the  acids  in  the  order  of  one  of  these  variables,  and  com- 
pare with  the  order  found  in  a  and  b. 

78.  Indicators. 

a.  Dilute  1  c.c.  of  concentrated  hydrochloric  acid  with  9  c.c. 
of  water,  mix  thoroughly  and  add  one  drop  of  this  diluted  acid 
to  100  c.c.  of  distilled  water  (A).  Mix  thoroughly  and  add 
equal  portions  to  four  test-tubes,  then  add  to  each  test-tube  a 


INDICATORS  87 

few  drops  of  one  of  the  indicators,  methyl  orange,  methyl  red, 
congo  red  and  litmus,  the  acid  and  alkali  colors  of  each  of  which 
should  be  previously  determined.  Dilute  50  c.c.  of  this  acid 
(A)  with  50  c.c.  of  distilled  water,  mix  thoroughly  and  test 
again  with  each  of  the  four  indicators,  and  repeat  this  diluting 
and  testing  until  each  of  the  four  indicators  have  changed  color. 
Which  indicator  changes  color  in  the  greatest  hydrogen  ion 
concentration?  which  in  the  least? 

b.  Make  up  a  solution  of  ammonium  hydroxide  in  exactly  the 
same  way  that  the  hydrochloric  acid  solution  was  prepared, 
using  the  concentrated  ammonium  hydroxide,  sp.  gr.  0.90, 
and  test  the  successive  dilutions  of  this  with  phenol phthalein, 
tropaolin  O  and  litmus.  Which  indicator  changes  color  in 
the  greatest  hydroxyl  ion  concentration?  which  in  the  least? 

c.  To  two  portions  of  10  c.c.  of  normal  sodium  acetate  add 
methyl  orange  and  phenolphthalein  respectively  as  indicators 
(is  the  reaction  of  sodium  acetate  acid  or  alkaline?)  then  titrate 
each  with  normal  acetic  acid  and  record  the  quantities  of  acid 
used  to  neutralize  the  solution,  with  each  indicator.     Explain 
why  these  quantities  shoud  be  different.     Now  repeat  the  pro- 
cedure titrating  with  normal  hydrochloric  acid  instead  of  acetic, 
record  the  quantities  used  and  explain  the  different  effects  of 
the  two  acids.     In  neutralizing  sodium  hydroxide  with  acetic 
acid  which  of  these  two  indicators  should  be  used?    In  neutral- 
izing with  hydrochloric  acid  would  it  make  much  difference 
which  was  used? 

d.  Test  a  solution  of  ammonium  sulfate  with  methyl  orange 
and  phenolphthalein,  and  titrate  equal  volumes  of  the  solution 
with  normal  ammonium  hydroxide  or  normal   sulfuric  acid, 
according  as  the  original  solution  reacts  acid  or  alkaline,  using 
these  two  indicators.     What  indicator  should  be  used  in  neu- 
tralizing a  strong  acid,  as  sulfuric  acid,  with  a  weak  base,  as 
ammonium  hydroxide?     Titrate  equal  quantities  of  normal 
ammonium  hydroxide  with  normal  acetic  acid,  using  methyl 
orange  and  phenolphthalein  as  indicators.     Will  these  indi- 
cators serve  to  show  the  neutral  point  in  this  case?     Why? 


88  METALS  AND  NON-METALS 

79.  Indicators.     Quantitative. 

a.  By  titrating  a  mixture  of  sodium  hydroxide  and  sodium 
bicarbonate  with  normal  hydrochloric  acid  it  is  possible  to 
titrate  only  the  sodium  hydroxide,  if  the  indicator  is  chosen  so 
that  as  soon  as  carbonic  acid  is  formed  it  shows  acid  reaction, 
and  with  another  indicator  which  is  not  affected  by  so  low  a 
hydrogen  ion  concentration  as  is  supplied  by  carbonic  acid, 
but  is  affected  as  soon  as  free  hydrochloric  acid  exists  in  the 
solution,  it  is  possible  to  determine  the  total  sodium  hydroxide 
plus  sodium  bicarbonate  in  the  solution.  Which  indicator 
should  be  used  for  each  titration?  Test  your  decision  by  try- 
ing a  solution  of  sodium  bicarbonate  alone  with  the  first  in- 
dicator, and  determine  how  much  normal  hydrochloric  acid  is 
required  to  give  acid  reaction,  and  by  trying  the  second  indi- 
cator to  determine  that  carbonic  acid  does  not  give  acid  reac- 
tion, but  that  free  hydrochloric  acid  does. 

6.  Titrate  equal  volumes  (pipette)  of  a  solution  of  these  two 
salts  with  normal  hydrochloric  acid,  using  the  two  indicators 
chosen,  and  from  these  two  titrations  calculate  the  amounts  of 
sodium  hydroxide  and  sodium  bicarbonate  present  in  the 
mixture.  In  the  original  solution  could  sodium  hydroxide  and 
sodium  bicarbonate  exist  together?  What  would  be  formed? 
Is  there  an  indicator  which  would  give  an  acid  reaction  with 
sodium  bicarbonate,  so  that  when  a  mixture  of  sodium  hydrox- 
ide and  sodium  carbonate  were  titrated  only  the  sodium  hydrox- 
ide would  be  neutralized?  Write  all  the  reactions  which  occur 
when  sodium  hydroxide  and  sodium  bicarbonate  are  mixed 
in  solution,  and  when  the  mixture  is  titrated,  first  with  one  indi- 
cator and  then  with  the  other.  From  the  results  of  the  titra- 
tion above  calculate  the  amount  of  sodium  carbonate  in  the 
given  mixture. 


CHAPTER  XV 
THE  ALKALI  METALS 

80.  The  Transition  of  Sodium  Sulfate  from  the  Anhydrous 
to  the  Hydrated  Salt. 

Add  30  grams  of  anhydrous  sodium  sulfate  (if  only 
the  hydrated  salt  is  available  calculate  the  quantity  re- 
quired to  give  30  grams  of  the  anhydrous  salt)  to  40  c.c.  of 
water  heated  to  50°  C.,  and  hold  the  mixture  at  that  tempera- 
ture for  about  five  minutes.  Is  the  solid  phase  the  hydrated 
or  the  anhydrous  salt?  Can  you  suggest  an  experimental  proof 
for  your  answer  to  this  question?  Now  cool  by  immersing  the 
beaker  in  water  at  15°  to  20°  C.  and  record  the  temperature  of  the 
solution  every  half  minute  for  ten  minutes,  stirring  it  constantly. 
If  the  temperature  falls  below  25°  drop  in  a  few  crystals  of  the 
hydrated  salt.  Make  a  plot  to  represent  the  change  in  tempera- 
ture with  the  time,  letting  the  horizontal  axis  represent  tem- 
perature and  the  vertical  axis  time,  and  explain  the  sudden 
break  in  the  curve.  Is  the  change  from  the  hydrated  to  the 
anhydrous  salt  endothermic  or  exothermic? 

81.  Sodium    Carbonate  by  the    Solvay  or  Ammonia  Process. 

(Adapted  from  Blanchard's  Synthetic  Inorganic  Chemistry.) 
a.  To  50  c.c.  of  concentrated  ammonium  hydroxide  (0.90 
specific  gravity)  add  150  c.c.  of  distilled  water.  Place  in  a 
flask  and  add  60  grams  of  sodium  chloride  free  from  lumps. 
Shake  until  the  salt  is  nearly  or  quite  dissolved,  and  filter  the 
solution,  if  it  is  not  perfectly  clear.  Pass  a  delivery  tube  through 
one  hole  of  a  two-hole,  tightly  fitting  stopper  placed  in  a  300 
c.c.  flask.  Provide  a  plug  for  the  other  hole.  Let  the  tube 
dip  into  the  solution  which  is  placed  in  the  flask,  and  pass  in 
carbon  dioxide  gas  from  a  Kipp  generator  until  all  the  air  has 

89 


90  THE  ALKALI  METALS 

been  displaced  from  the  flask;  then  close  the  flask  and  allow  the 
gas  to  pass  in  as  fast  as  it  will  be  absorbed.  What  is 
formed?  Occasionally  as  the  action  seems  to  slacken,  loosen 
the  plug  for  a  moment.  Shake  the  flask  frequently.  It  will 
take  several  hours  for  the  solution  to  absorb  sufficient  carbon 
dioxide,  and  it  may  be  left  over  night  connected  with  the 
generator.  When  no  more  gas  can  be  absorbed  pour  the 
mixture  from  the  flask  onto  a  fluted  filter  paper,  wash 
the  precipitate  with  a  small  amount  of  cold  water,  and 
allow  it  to  dry  in  the  air.  Look  up  the  solubilities  of  the 
substances  started  with  and  the  final  products  of  this 
reaction  and  explain  upon  what  the  success  of  the  process 
depends.  Why  could  not  ammonium  chloride  be  used  instead 
of  ammonium  hydroxide?  What  remains  in  the  filtrate  from 
the  bicarbonate?  Would  the  process  be  financially  a  success 
if  this  filtrate  was  discarded  ?  How  may  the  ammonium  hydrox- 
ide be  regenerated  from  it  and  what  is  the  final  waste  product  ? 

b.  Add  hydrochloric  acid  to  some  of  the  solid  bicarbonate 
prepared  above.     Heat  a  small  portion  and  test  the  gas  evolved 
for  carbon  dioxide.     To  the  residue  after  heating  add  hydro- 
chloric acid.     What  is  the  reaction  on  heating? 

c.  Heat  about  1  gram  of  the  bicarbonate  till  decomposition 
is  complete,  dissolve  the  product  in  very  little  water,  test  with 
litmus  and  compare  with  the  effect  of  sodium  bicarbonate 
solution  on  litmus.     Add  a  solution  of  lime  water  (calcium 
hydroxide)  till  there  is  no  further  precipitation,  and  filter  the 
precipitate?     What  is  in  solution? 

82.  Potassium  Nitrate  from  Sodium  Nitrate  and  Potassium 
Chloride.  (Adapted  from  Blanchard's  Synthetic  Inorganic 
Chemistry.) 

Dissolve  50  grams  of  sodium  nitrate  and  44  grams  of  potas- 
sium chloride  in  100  c.c.  of  water  and  evaporate  on  the  water 
bath  in  a  porcelain  dish,  adding  the  solution  in  portions  if  the 
dish  will  not  hold  it  all  at  once,  till  the  solution  is  half  its 
original  volume.  Without  letting  the  liquor  cool  separate  it 


POTASSIUM  NITRATE  91 

from  the  crystals  which  have  formed  during  the  evaporation. 
This  may  be  done  by  decantation  or  by  filtering  through  a  fluted 
filter  paper  which  has  been  wet  with  hot  water.  Allow  the 
filtrate  to  cool  by  holding  the  beaker  containing  it  under  run- 
ning water,  and  filter  out  the  second  crop  of  crystals.  What 
four  salts  may  be  present  in  the  solution  after  potassium  chlo- 
ride and  sodium  nitrate  have  been  mixed  together?  Look  up 
the  solubilities  of  each  of  these  salts  and  determine  which  would 
separate  out  first  from  a  solution  containing  equi-molecular 
amounts  of  each,  on  evaporating  at  100°.  Test  the  first  crop 
of  crystals  to  determine  if  your  reasoning  is  correct.  If  half 
the  sodium  and  chloride  ions  in  the  original  mixture  were 
removed,  what  would  be  left  in  the  solution  and  in  what  pro- 
portions? Considering  their  solubilities  at  high  and  low  tem- 
perature, and  the  relative  amounts  present,  which  salt  would 
you  expect  would  separate  out  on  cooling  the  solution?  Devise 
and  apply  a  test  on  the  second  crop  of  crystals  to  determine 
whether  your  conclusion  is  correct.  Why  is  potassium  nitrate 
used  for  the  manufacture  of  gunpowder  rather  than  sodium 
nitrate? 


CHAPTER  XVI 
COPPER  AND  SILVER 

83.  Copper  and  Cupric  Compounds. 

a.  Treat  bright  copper  filings  with  dilute  sulfuric  and  dilute 
hydrochloric  acids.  Does  the  metal  dissolve  in  either  case? 
Pour  off  the  dilute  acids  and  replace  them  with  concentrated 
acids,  warm,  and  note  the  effect  in  each  case.  Test  the  effect 
of  both  dilute  and  concentrated  nitric  acid  on  copper  and  ac- 
count for  the  effect  of  each  acid  on  the  copper.  Will  copper 
displace  hydrogen  from  an  acid? 

6.  Heat  some  copper  filings  in  an  open  crucible  and  note  the 
change  in  appearance.  Test  the  product  for  solubility  in  dilute 
sulfuric  and  dilute  hydrochloric  acids,  and  account  for  the  dif- 
ference in  behavior  of  copper  before  and  after  heating  in  air. 

c.  Add  ammonium  hydroxide  drop  by  drop  to  about  5  c.c.  of 
copper  sulfate  solution,  shaking  after  each  addition,  until  the 
mixture  smells  strongly  of  ammonia.  Account  for  every  change 
which  occurs.  To  another  portion  of  copper  sulfate  solution 
add  sodium  hydroxide  solution  till  the  mixture  is  alkaline,  then 
heat,  and  account  for  every  change  which  occurs. 

84.  Cuprous  Compounds. 

a.  Prepare  a  solution  of  3.5  grams  of  copper  sulfate  in  50  c.c. 
of  water,  and  another  solution  of  17  grams  of  Rochelle  salt 
(sodium  potassium  tartrate)  and  7  grams  of  potassium  hydrox- 
ide (or  its  equivalent  of  sodium  hydroxide)  in  50  c.c.  of  water. 
Mix  these  solutions  and  heat  to  boiling  then  add  a  solution  of 
glucose  until  all  the  copper  is  precipitated.  How  can  that 
point  be  recognized?  Filter  and  wash  the  precipitate  (what  is 
it?)  with  warm  water  several  times,  then  allow  to  dry. 

6.  To  a  small  portion  of  the  precipitate  add  dilute  hydro- 
chloric acid.  What  is  formed?  Test  different  portions  of  this 

92 


CUPROUS  COMPOUNDS  93 

white  precipitate  for  solubility  in  concentrated  hydrochloric 
acid  and  in  ammonium  hydroxide.  Is  the  cupric  ion  present 
in  either  of  these  solutions?  State  the  reason  for  your  answer. 
Pour  the  concentrated  hydrochloric  acid  solution  into  an  excess 
of  water  and  explain. 

c.  Add  ammonium  hydroxide  to  a  small  portion  of  the  red 
cuprous  oxide  and  explain  the  color  changes. 

d.  Add  dilute  sulfuric  acid  to  another  portion  of  the  cuprous 
oxide.     Note  the  color  of  the  solution  and  of  the  insoluble 
residue.     Devise  and  apply  tests  to  prove  that  this  residue  is  not 
cuprous  oxide,  but  is  metallic  copper.     Is  the  cuprous  or  cupric 
ion  present  in  the  solution?     What  is  the  oxidizing  and  what  the 
reducing  agent  in  this  reaction? 

e.  To  a  few  cubic  centimeters  of  copper  sulfate  solution  add 
an  excess  of  potassium  iodide  solution,  filter  and  wash  the 
precipitate  first  with  a  little  potassium  iodide  solution  (why?), 
then  with  water.     Explain  the  color  of  the  filtrate.     To  one 
portion  of  the  precipitate  add  dilute  potassium  hydroxide  and 
warm.     Is  this  the  oxide  of  cupric  or  cuprous  copper?     To 
another  portion  of  the  precipitate  add  ammonium  hydroxide. 
What  is  formed? 

85.  Electromotive  Series. 

Place  the  following  mixtures  in  test-tubes,  taking  especial 
pains  to  see  that  the  metal  has  a  bright,  shiny  surface,  i.e.,  is 
not  covered  with  an  oxide.  Observe  the  changes  occurring, 
and  complete  the  chemical  equations 

Cu+ZnS04  solution.  Pb-f-ZnS04  solution. 

Cu  +  Pb(N03)2  solution.         Pb  +  Cu(N03)2  solution. 

Cu+Eg(N08)j  solution.         Pb  +  Hg(N03)2  solution. 

Hg+ZnS04  solution. 

Hg  +  Pb(N03)2  solution. 

Hg  +  CuS04  solution. 

Zn  +  Pb(N03)2  solution.         Pb  +  HCl  (dilute). 

Zn+Cu(N08)2  solution. 

Zn  +  Hg(N03)2  solution. 


94  COPPER  AND  SILVER 

These  reactions  are  not  instantaneous  and  should  be  allowed  to 
continue  during  the  time  of  at  least  one  laboratory  period. 
Arrange  the  four  elements  in  the  order  of  their  "  solution 
tension"  or  tendency  to  change  from  the  free  to  the  ionic  con- 
dition. Refer  to  the  last  part  of  Exp.  13  and  find  the  place 
for  hydrogen  in  the  series. 

86.  Silver  Oxide. 

a.  Prepare  silver  oxide  by  adding  an  excess  of  potassium  hy- 
droxide to  about  20  c.c.  of  silver  nitrate  solution,  filter  out  the 
oxide  and  wash,  then  test  small  portions  of  the  precipitate  for 
solubility   in  ammonium  hydroxide  and  dilute  sulfuric  acid. 
Heat  the  remainder  of  the  precipitate  in  a  porcelain  crucible, 
and  note  the  change  in  appearance  and  test  small  portions 
for  solubility  in  ammonium  hydroxide  and  dilute  sulfuric  acid. 
Explain  the  difference  in  conduct  before  and  after  heating. 
Dissolve  the  remaining  silver  in  concentrated  nitric  acid  which  has 
been  diluted  with  an  equal  volume  of  water,  avoiding  an  excess 
of  the  acid,  then  evaporate  the  solution  till  crystals  begin  to 
form  on  the  surface,  then  crystallize  by  cooling  and  spontaneous 
evaporation. 

b.  Convert  a  small  portion  of  the  silver  nitrate  into  silver 
chloride,  add  ammonium  hydroxide  and  explain.     Devise  and 
apply  a  method  for  converting  the  rest  to  silver  sulfate.     Com- 
pare the  properties  of  the  cuprous  and  silver  ions. 

87.  Silver  Chloride  Emulsion. 

Prepare  a  1  per  cent  gelatin  solution  (gelatin  will  dissolve 
readily  on  warming  with  water)  and  to  a  measured  amount  of 
the  solution  (about  25  c.c.)  add  two  drops  of  concentrated 
hydrochloric  acid  and  an  excess  of  silver  nitrate  solution  and 
compare  with  the  appearance  of  the  silver  chloride  when  the 
same  procedure  is  followed,  using  distilled  water  instead  of  the 
gelatin  solution.  Allow  the  two  preparations  to  stand  for  an 
hour  or  more  and  again  compare.  What  is  the  technical  impor- 
tance of  silver  chloride  emulsions? 


CHAPTER  XVII 
CALCIUM,  STRONTIUM  AND  BARIUM 

88.  Calcium  Carbonate  and  Lime. 

a.  Heat  about  2  grams  of  well  ground  marble  in  a  hard  glass 
test  -tube  fitted  with  a  stopper  and  delivery  tube,  and  test  the 
gas  evolved  by  passing  it  through  lime  water.  Remove  small 
portions  of  the  material  in  the  test-tube  at  intervals  of  about  fif- 
teen minutes,  and  test  with  hydrochloric  acid  to  determine  the 
rate  of  decomposition.  How  could  marble  be  kept  from  decom- 
posing at  800°  ?  After  the  heating  has  been  continued  for  about 
an  hour  test  a  portion  of  the  residue  in  the  test-tube  for  solubility 
in  water  and  for  the  reaction  of  the  solution  toward  litmus,  and 
compare  with  the  original  marble.  Account  for  the  differences 
noted.  Filter  out  the  insoluble'  residue  and  pass  carbon  dioxide 
into  the  filtrate.  What  is  formed?  Add  a  little  quicklime  to 
water  and  note  the  temperature  change,  and  the  reaction  of 
the  solution  toward  litmus. 

89.  The  Relative  Solubility  of  the  Hydroxides  of  Calcium,  Stron- 
tium and  Barium.    Quantitative. 

a.  Make  saturated  solutions  of  barium,  strontium  and  calcium 
hydroxides  by  adding  an  excess  of  the  solid  to  50  c.c.  of  distilled 
water  and  allowing  to  stand  in  stoppered  flasks  until  the  super- 
natant liquid  is  clear.  (Keep  these  solutions  for  Exp.  90.) 
Now  make  200  c.c.  of  an  approximately  0.1  normal  hydro- 
chloric acid  solution  by  diluting  the  concentrated  acid  (sp.  gr. 
1.16).  Take  10  c.c.  of  each  of  the  hydroxide  solutions,  measured 
with  a  pipette,  from  the  clear  supernatant  liquid.  Neutralize 
these  with  the  acid  from  a  burette,  recording  the  exact  quantity 
of  acid  required  in  each  case  and  using  an  indicator  which  will 
not  give  the  acid  color  in  so  slight  a  hydrogen  ion  concentration 

95 


96  CALCIUM,  STRONTIUM  AND  BARIUM 

as  comes  from  carbonic  acid.  Why?  Calculate  the  relative 
solubilities  of  these  three  hydroxides.  If  the  acid  is  assumed  to 
be  exactly  0.1  normal  what  are  the  absolute  solubilities? 

b.  Titrate  100  c.c.of  tap  water  in  the  same  way.  What  is  the 
alkaline  substance?  Assuming  that  the  alkali  in  the  tap  water 
is  lime,  express  the  concentration  of  the  tap  water  in  terms  of  a 
saturated  solution  of  lime.  (That  is,  if  a  saturated  solution 
was  taken  as  unity,  this  would  be  some  fraction  of  a  saturated 
solution.) 

90.  The  Relative  Solubilities  of  Calcium,  Strontium  and  Barium 
Bicarbonates. 

Measure  out  10  c.c.  of  each  of  the  saturated  solutions  of  cal- 
cium, strontium  and  barium  hydroxides  prepared  in  Exp.  89 
and  pass  carbon  dioxide  from  a  Kipp  generator  into  each  of  these 
solutions,  adding  water  to  each  gradually,  while  the  carbon 
dioxide  continues  to  pass  through,  until  the  precipitate  first 
formed  just  dissolves  in  the  excess  of  water  saturated  with 
carbon  dioxide.  (A  small  granular  precipitate  in  the  case  of  the 
barium  may  be  ignored.)  Now  measure  the  total  volumes  of 
each  of  the  solutions,  and  calculate  the  relative  solubilities  of 
the  bicarbonates  from  that  of  the  hydroxides  (Exp.  89)  and 
these  measured  volumes. 

91.  The  Relative  Solubilities  of  Calcium,  Strontium  and  Barium 
Sulfates. 

Prepare  calcium  sulfate  from  a  small  piece  of  marble  (can 
this  be  done  by  adding  sulfuric  acid  directly?  Why?).  Filter 
and  wash  the  calcium  sulfate  free  from  sulfuric  acid  (would  you 
test  this  with  litmus  or  with  barium  chloride?  Why?)  then 
make  a  saturated  solution  of  this  calcium  sulfate  and  add  the 
clear  solution  to  strontium  chloride  solution  and  warm.  What  is 
precipitated?  What  remains  in  solution?  Add  the  filtrate 
from  this  to  barium  chloride  solution  and  explain.  Draw  con- 
clusions regarding  the  relative  solubility  of  barium,  strontium 
and  calcium  sulfates.  Explain  why  calcium  sulfate  precipi- 
tates from  solutions  of  calcium  chloride  and  sulfuric  acid  more 


CALCIUM  CHLORIDE  97 

slowly  than  barium  sulfate  from  solutions  of  barium  chloride 
and  sulfuric  acid. 

92.  Preparation  and  Properties  of  Calcium  Chloride. 

a.  Prepare  a  solution  of  calcium  chloride  from  about   25 
grams  of  marble,  taking  care  to  avoid  an  excess  of  hydrochloric 
acid,  filter  and  crystallize  by  boiling  until  crystals  begin   to 
separate,  then  allow  to  cool,  remove  the  crystals  from  the 
mother  liquor,  and  dry  them  on  filter  paper.     Boil  down  the 
mother  liquor  and  continue  heating,  with  stirring,  until  the 
solid  residue  has  become  quite  white  and  opaque.     What  is  the 
difference  between  these  two  preparations?     How  much  of  the 
crystallized    salt    would    be    equivalent    to   a    gram   of    the 
anhydrous  ? 

b.  Add  equivalent  quantities  of  each  of  these  salts  (finely 
ground)  to  10  c.c.  of  water,  taking  the  temperature  before  addi- 
tion and  after  solution  of  the  salt,  and  account  for  the  difference. 

c.  Repeat  b,  using  equal  weights  of  finely  chopped  ice  or 
snow  instead  of  water,  and  account  for  the  difference  in  effect 
of  the  two  salts,  and  for  the  difference  in  effect  as  compared  with 
the  previous  test  with  water.     Allow  samples  of  each  of  the 
preparations  to  stand  in  the  open  air  and  account  for  the  change. 


CHAPTER  XVIII 
MAGNESIUM  ZINC  AND  MERCURY 

93.  The  Equivalent  Weight  of  Magnesium.    Quantitative. 

Prepare  an  apparatus  as  shown  in  Fig.  20,  in  which  an  ordi- 
nary test-tube  is  fitted  with  a  two  hole  stopper,  one  hole  of 


FIG.  20 


which  carries  a  glass  rod  which  has  been  bent  over  just  a  little 
at  its  lower  end,  and  after  being  put  into  the  cork  has  been  bent 

98 


ZINC  99 

at  right  angles  at  the  other  end,  while  the  other  hole  carries  a 
tube  bent  at  an  acute  angle,  to  connect  with  the  gas  burette. 
Make  the  apparatus  air  tight,  then  fill  the  burette  to  the  zero 
mark  with  water,  place  about  10  c.c.  of  dilute  hydrochloric 
acid  in  the  test-tube,  and  hang  on  the  glass  rod  a  weighed 
quantity  of  magnesium  ribbon  (not  more  than  0.05  gram)  or  a 
measured  length,  the  weight  of  which  can  be  calculated  from  a 
factor,  weight  per  centimeter,  on  the  bottle.  Connect  the  ap- 
paratus, open  the  burette  cock  and  read  the  volume  (at  atmos- 
pheric pressure),  drop  the  magnesium  ribbon  into  the  acid  by 
turning  the  glass  rod,  and  after  the  magnesium  is  all  dissolved 
read  the  volume  again,  noting  the  temperature  of  the  room  and 
the  atmospheric  pressure.  Calculate  the  weight  of  hydrogen 
evolved,  and  the  equivalent  weight  of  magnesium. 

94.  Magnesium  Carbonate  and  Oxide. 

Heat  magnesium  carbonate  as  calcium  carbonate  was  heated 
(Exp.  88),  testing  a  small  portion  with  hydrochloric  acid  every 
ten  minutes  till  decomposition  is  complete,  add  water  and  test 
the  suspension  with  litmus.  Compare  these  with  the  same 
experiments  on  calcium.  Add  a  saturated  solution  of  am- 
monium chloride  to  the  suspension  of  magnesium  oxide,  test 
again  with  litmus  and  account  for  the  changes. 

95.  The  Properties  of  Metallic  Zinc. 

Heat  a  granule  of  zinc  in  an  open  crucible,  noting  and  explain- 
ing all  changes  which  take  place  during  the  heating  and  subse- 
quent cooling.  Compare  with  the  conduct  of  silver  under  sim- 
ilar conditions.  Tabulate  the  melting-points  of  the  metals  of 
the  first  and  second  groups  of  the  periodic  system  and  formulate 
a  statement  of  the  relation  of  melting-points  to  atomic  weights. 

96.  The  Equivalent  Weight  of  Zinc.    Quantitative. 

Repeat  the  manipulation  described  in  Exp.  93  using  about 
0.1  gram  of  zinc,  and  place  in  the  test-tube  a  small  piece  of 
platinum  wire  with  which  the  zinc  may  come  in  contact,  to 
hasten  its  solution.  Calculate  the  equivalent  weight  of  zinc. 


100  MAGNESIUM,  ZINC  AND  MERCURY 

97.  Preparation  of  Zinc  Sulfate. 

a.  Make  a  suspension  of  5  grams  of  finely  powdered  zinc 
oxide  in  water,  test  this  with  litmus,  and  dissolve  by  adding 
acetic  acid,  or  dissolve  in  dilute  hydrochloric  acid  and  add 
sodium  acetate.  Explain  why  these  two  procedures  would 
give  the  same  hydrogen  ion  concentration.  Place  this  solution 
in  an  erlenmeyer  flask  fitted  with  a  one  hole  stopper  carrying  a 
delivery  tube  projecting  below  the  surface  of  the  liquid,  and 
pass  in  hydrogen  sulfide  with  the  stopper  loose  until  the  air  in 
the  flask  has  been  displaced,  then  close  the  flask,  when  hydrogen 
sulfide  will  be  supplied  as  fast  as  absorbed,  and  allow  the  flask 
to  stand  until  the  solution  is  saturated.  What  is  the  precipi- 
tate? Filter  and  wash,  then  dry  the  moist  precipitate  in  an 
evaporating  dish,  stirring  constantly,  and  when  entirely  dry 
heat  in  an  open  crucible  to  dull  red  for  half  an  hour,  or  until 
a  small  test  portion  shows  the  presence  of  appreciable  amounts 
of  a  soluble  sulfate.  Account  for  the  formation  of  a  sulfate. 
Write  the  equation.  Dissolve  in  water,  warm  and  filter.  Test 
the  filtrate  with  litmus  and  explain,  then  evaporate  the  solution 
to  crystallization.  Devise  and  apply  a  test  to  determine 
whether  the  insoluble  residue  is  mostly  zinc  oxide  or  zinc  sulfide. 

6.  Dissolve  the  zinc  sulfate  prepared  in  a,  acidify  with  dilute 
sulfuric  acid  and  pass  hydrogen  sulfide  into  a  portion  of  the 
solution,  then  add  an  equal  volume  of  sodium  acetate  solution 
and  again  pass  in  hydrogen  sulfide.  Account  for  the  difference. 

c.  To  two  other  portions  of  the  zinc  sulfate  solution  add 
sodium  hydroxide  till  a  permanent  precipitate  forms,  then  to  one 
add  an  excess  of  sodium  hydroxide,  and  to  the  other  hydro- 
chloric or  sulfuric  acid.  Is  zinc  hydroxide  a  base  or  an  acid? 
Indicate  its  products  of  ionization  and  the  effect  which  the 
presence  of  a  base  or  an  acid  has  upon  its  ionization. 

98.  The  Preparation  and  Properties  of  Mercurous  and  Mercuric 
Nitrate. 

a.  Clean  5  grams  of  mercury  by  pouring  through  a  dry  filter 
paper  through  which  a  pin  hole  has  been  pricked  at  the  apex. 


MERCUROUS  AND  MERCURIC  NITRATES       101 

Calculate  the  amount  of  concentrated  nitric  acid  required  to 
convert  this  to  mercurous  nitrate  (HgN03)  and  take  a  little  less 
(why?)  than  the  calculated  amount  of  the  concentrated  acid 
(specific  gravity  1.40,)  and  dilute  with  twice  its  volume  of  water. 
Dissolve  the  mercury  in  this,  warning  gen.tly  whe,n,  necessary 
to  keep  the  reaction  going,  and  after. -it 'has 'apparently  ceased 
(a  good  indication  of  complete  reaction  is:tJ¥6',fd7it^tioAtof  a 
yellow  precipitate  on  diluting  a  small  portion'  of  the  solution' with 
an  equal  volume  of  water — explain  this  test  after  the  experi- 
ment has  been  completed)  pour  off  the  solution  from  the  residue 
of  undissolved  mercury,  cool  and  allow  to  crystallize.  To 
separate  portions  of  the  mother  liquor  add  hydrochloric  and 
sulfuric  acids  and  ammonium  hydroxide.  Compare  the  solu- 
bility of  mercurous  nitrate,  sulfate  and  chloride.  Add  water 
to  some  of  the  crystals,  then  warm  and  test  the  water  solution 
with  litmus.  Account  for  the  test.  What  is  the  precipitate? 
Is  it  soluble  in  water?  In  nitric  acid? 

b.  Dissolve  the  rest  of  the  mercurous  nitrate  in  as  little  dilute 
nitric  acid  as  possible  and  warm,  adding  concentrated  nitric  acid 
as  necessary  till  brown  fumes  are  no  longer  given  off.  What 
reaction  is  taking  place  ?  Calculate  the  amount  of  concentrated 
nitric  acid  required  to  convert  5  grams  of  mercury  to  mercuric 
nitrate  (Hg(N03)2)  and  compare  with  the  quantity  required  for 
mercurous  nitrate.  Crystallize  this  salt.  How  does  its  solu- 
bility compare  with  that  of  mercurous  nitrate?  Dissolve  some 
of  the  crystals  in  water,  adding  dilute  nitric  acid  if  necessary, 
and  test  the  solution  with  hydrochloric  and  sulfuric  acids  and 
ammonium  hydroxide,  and  compare  with  the  mercurous  com- 
pounds. Does  the  ammonium  hydroxide  form  mercuric  hydrox- 
ide? Devise  and  apply  a  test  to  determine  this.  Add  some 
bright  metallic  copper  to  the  mercuric  nitrate  solution.  Where 
is  mercury  in  the  electromotive  series?  What  metals  will  be 
precipitated  from  solution  by  mercury?  Add  a  small  amount 
of  zinc  to  mercuric  nitrate  solution  and  warm  till  the  zinc  is 
nearly  all  dissolved,  then  test  the  solution  by  adding  hydro- 
chloric acid.  What  was  the  effect  of  the  zinc? 


CHAPTER  XIX 
—    -'ALUMINIUM 

99.  Aluininates  and  Aluminium  Hydroxide. 

Dissolve  about  2  grams  of  aluminium  in  dilute  potassium 
hydroxide,  avoiding  an  excess  of  the  alkali.  What  remains  in 
solution?  What  gas  is  evolved?  Filter  the  solution  and  pass 
hydrogen  sulfide  into  a  small  portion.  Is  a  sulfide  precipitated? 
Prove  your  answer  experimentally.  Pass  carbon  dioxide  into 
the  main  portion  of  the  solution.  What  is  the  precipitate? 
Filter  and  wash,  and  show  experimentally  that  it  is  not  a  car- 
bonate. Test  small  portions  of  this  precipitate  for  solubility  in 
acids  and  alkalis,  and  explain  in  terms  of  the  ionization  theory. 
Dissolve  the  main  portion  of  the  precipitate  in  hydrochloric 
acid,  evaporate  until  crystals  appear  on  the  surface  of  the 
liquid,  which  should  be  kept  strongly  acid,  and  crystallize  by 
cooling  and  spontaneous  evaporation.  Dry  the  crystals,  dis- 
solve a  small  portion  in  water  and  pass  in  hydrogen  sulfide. 
Compare  with  the  action  of  hydrogen  sulfide  on  the  alkaline 
solution  and  explain.  Heat  the  rest  of  the  crystals  in  a  dry 
test-tube  till  decomposition  is  complete,  noting  the  odor  of  the 
gas  given  off.  What  is  it?  Is  the  residue  a  chloride?  Test 
this  experimentally. 

100.  Water  of  Crystallization  in  Alum.    Quantitative. 

Calculate  the  quantities  of  aluminium  sulfate  (hydrated)  and 
potassium  sulfate  necessary  to  make  5  grams  of  the  crystallized 
aluminium  alum,  and  dissolve  these  amounts  in  20  c.c.  warm 
water,  cool  and  allow  to  crystallize  slowly.  Dry  the  crystals  on 
filter  paper,  then  by  standing  in  the  open  air,  and  after  they  are 
thoroughly  dry  determine  the  water  of  crystallization  by 
placing  about  1  gram  in  a  clean,  dry  crucible  which  has  been 

102 


ALUM  103 

previously  weighed,  and  weighing  accurately.  Cover  this  crucible 
and  place  in  another  crucible  of  the  same  size  and  heat  with  a 
flame  not  more  than  two  inches  high,  and  at  least  two  inches  from 
the  bottom  of  the  lower  crucible.  After  most  of  the  water  has 
been  driven  off  the  lower  crucible  may  be  removed  and  the  small 
flame  brought  just  below,  but  not  quite  touching  the  crucible 
containing  the  alum,  and  the  heating  continued  until  constant 
weight  is  attained.  Too  high  a  temperature  will  cause  decom- 
position of  the  aluminium  sulfate.  Calculate  the  per  cent  of 
water  in  alum  from  this  experiment,  and  from  the  formula,  and 
compare. 
101.  Aluminium  in  Kaolin. 

Fuse  some  potassium  pyrosulfate  in  a  crucible  and  add  pow- 
dered kaolin  in  small  portions  until  the  mass  will  no  longer 
melt,  then  cool,  dissolve  in  boiling  water,  filter  and  add  am- 
monium hydroxide  to  the  filtrate.  What  is  formed?  Test 
its  solubility  in  potassium  hydroxide  and  in  hydrochloric 
acid. 


CHAPTER  XX 
LEAD  AND  TIN 

102.  The  Properties  of  Metallic  Lead  and  of  Lead  Salts. 

a.  Test  the  solubility  of  metallic  lead  in  hydrochloric,  acetic, 
sulfuric  and  nitric,  acids.     To  solutions  of  lead  nitrate  add  each 
of  the  above  acids,  first  cold,  then  warm.     From  the  informa- 
tion so  gained  explain  the  effects  of  the  acids  on  metallic  lead. 
From  the  position  of  lead  in  the  electromotive  series  would  you 
expect  it  to  dissolve  in  these  acids? 

b.  Add  sodium  hydroxide  solution  to  lead  nitrate  solution, 
first  in  small  amount  then  in  excess,  and  pass  hydrogen  sulfide 
into  the  alkaline  solution.     Devise  and  apply  a  test  to  determine 
whether  the   sulfide  or  hydroxide  is  precipitated  here.     Pass 
hydrogen  sulfide  into  a  hot  solution  of  lead  acetate,  acidified 
with  hydrochloric  acid.     Is  the  sulfide  or  hydroxide  precipi- 
tated?    Compare  these  results  with  those  obtained  on  similar 
treatment  of  aluminium.     Which    element   shows    the  more 
basic  properties?    From  the  positions  of  lead  and  aluminium 
in  the  Periodic  System  which  would  you  expect  to  show  most 
basic  properties? 

103.  The  Oxides  of  Lead. 

a.  Heat  a  small  sample  of  red  lead  in  a  dry  test-tube,  noting 
the  change  in  appearance,  and  at  the  same  time  test  for  oxygen. 
After  decomposition  is  complete  divide  the  residue  into  two 
portions;  to  one  add  dilute  nitric  acid  and  to  the  other  sodium 
hydroxide  solution.  Shake  and  filter  if  necessary,  then  pass 
hydrogen  sulfide  into  each  solution.  What  has  been  the  re- 
action in  each  case?  Is  each  color  change  on  heating  the  red 
lead  accompanied  by  decomposition?  Test  this  by  heating 
another  small  portion  till  only  the  first  change  has  occurred, 
then  cooling. 

6.  Treat  another  small  sample  of  red  lead  with  dilute  nitric 
acid,  stirring  till  the  change  is  complete.  Filter  and  add  sodium 
hydroxide  solution  to  the  filtrate,  first  in  small  amounts  then 

104 


TIN  105 

in  excess,  and  account  for  the  changes.  What  is  the  residue 
insoluble  in  nitric  acid?  Is  it  soluble  in  sodium  hydroxide 
solution?  In  hydrochloric  acid  (hood)?  What  are  the  prod- 
ucts of  the  reaction  with  hydrochloric  acid?  Warm  to  dis- 
solve the  ]ead  chloride  and  test  the  solution  by  adding  sodium 
Itydroxids  solution  and  compare  with  the  reaction  of  the  nitrate 
(the  first  filtrate,  above). 
104.  Stannous  and  Stannic  Tin. 

a.  Dissolve  about  1  gram  of  tin  in  concentrated  nitric  acid 
diluted  with  an  equal  volume  of  water,  filter  and  wash.  What 
is  the  precipitate?  Keep  the  filtrate  and  at  the  end  of  this 
experiment  devise  and  apply  a  test  to  determine  the  presence 
or  absence  of  tin.  Dissolve  a  small  portion  of  the  precipitate 
in  potassium  hydroxide  solution.  Dissolve  the  rest  in  as  little 
dilute  hydrochloric  acid  as  possible  (A).  What  is  formed  in 
each  case?  Test  small  portions  of  each  of  these  solutions  by 
saturating  with  hydrogen  sulfide,  then  acidify  with  hydrochloric 
acid  and  account  for  every  change  that  takes  place  in  each  solu- 
tion. Dilute  a  small  portion  of  the  hydrochloric  acid  solution 
(A)  and  add  potassium  permanganate  solution  (hydrochloric 
acid  may  reduce  potassium  permanganate  solution  if  sufficiently 
concentrated).  Is  this  stannous  or  stannic  tin? 

6.  Dissolve  1  gram  of  tin  by  warming  with  concentrated 
hydrochloric  acid  diluted  with  an  equal  volume  of  water,  avoid- 
ing an  excess  of  the  acid.  Dilute  and  add  potassium  perman- 
ganate solution  to  a  small  portion.  Account  for  the  effect. 
Add  potassium  hydroxide  solution  to  another  portion  till  a 
permanent  precipitate  forms,  which  on  addition  of  more  potas- 
sium hydroxide  and  warming  redissolves.  What  is  in  solution? 
Pass  hydrogen  sulfide  into  this  alkaline  solution  and  also  into 
the  acid  solution.  Compare  with  the  same  tests  in  a  and  ac- 
count for  the  difference. 

105.  The  Hydrogen  Equivalent  of  Tin  and  Its  Atomic  Weight. 
Quantitative. 

a.  Repeat  the  manipulation  described  in  Exp.  93  using  about 
0.1  to  0.15  gram  of  tin,  and  place  in  the  test-tube  a  small  piece 


106  LEAD  AND  TIN 

of  platinum  wire  with  which  the  tin  may  come  in  contact,  to 
hasten  its  solution.  Dissolve  the  tin  in  concentrated  hydro- 
chloric acid  instead  of  the  dilute,  warming  if  necessary. 
Calculate  the  equivalent  weight  of  tin,  and  by  referring  to 
Exp.  52  calculate  the  equivalent  weight  of  tin  from  that 
experiment.  Account  for  the  difference. 

b.  Apply  Dulong  and  Petit's  Law  to  determine  the  atomic 
weight  from  these  equivalents  by  determining  the  specific  heat 
as  follows.  Take  a  piece  of  compact  tin  weighing  5  to  10  grams, 
or  if  a  compact  piece  is  not  available  melt  the  required  amount 
of  feathered  or  granulated  tin  in  a  crucible,  and  file  down  until 
a  bright  smooth  surface  is  obtained.  Weigh  this  to  within 
0.1  gram,  place  in  a  dry  test-tube,  and  immerse  the  test-tube  in 
a  beaker  of  boiling  water  for  ten  to  fifteen  minutes,  until  the  tin 
shall  have  reached  the  temperature  of  the  water.  Measure  out 
in  a  small  dry  beaker  25  c.c.  of  water  at  the  room  temperature, 
and  provide  it  with  a  stirrer  made  by  flattening  the  end  of  a 
glass  rod.  Read  the  temperature  of  the  water  carefully, 
remove  the  tin  from  the  boiling  water  and  drop  it  directly 
from  the  test-tube  into  the  beaker  of  water.  Stir  well,  and  after 
a  moment  or  two  read  the  temperature  of  the  water  again.  If 
the  temperature  is  read  too  soon  error  will  result  from  incom- 
plete distribution  of  the  heat  between  the  tin  and  the  water, 
if  too  late,  there  may  be  an  appreciable  error  introduced  by 
loss  of  heat  by  radiation.  Now  read  the  temperature  of  the 
boiling  water  and  from  this  data  calculate  the  specific  heat  of 
the  tin  as  follows:  The  change  in  temperature  of  the  water 
times  its  volume  measured  in  c.c.  gives  roughly  the  number 
of  calories  supplied  by  the  tin  (this  is  introducing  a  slight 
error  by  ignoring  the  heat  capacity  of  the  beaker  and  stirrer, 
and  loss  of  heat  by  radiation),  and  that  divided  by  the  change 
in  temperature  of  the  tin  times  its  weight  gives  the  amount  of 
heat  supplied  by  1  gram  of  the  tin  in  changing  its  tempera- 
ture by  1°,  which  is  its  specific  heat.  Apply  Dulong  and 
Petit's  Law  and  give  the  true  atomic  weight  of  tin. 


CHAPTER  XXI 
CHROMIUM 

106.  The  Preparation  of  Sodium  Chromate  from  Chrome  Iron 
Ore. 

Mix  thoroughly  5  grams  of  sodium  peroxide  with  2  grams  of 
well  ground  chrome  iron  ore  and  heat  gently,  keeping  the 
mixture  just  melted  for  about  twenty  minutes.  Cool  and  dis- 
solve in  a  small  amount  of  warm  water  and  filter.  What  is  the 
precipitate?  Add  hydrochloric  acid  to  the  filtrate  as  long  as  the 
orange  color  formed  as  the  acid  comes  in  contact  with  the 
solution  changes  back  to  the  original  yellow  color  on  stirring. 
Explain  these  changes  after  the  experiment  is  completed. 
What  two  salts  are  now  present  in  this  solution?  Look  up 
their  solubilities  and  devise  a  method  of  taking  advantage  of 
their  differences  in  solubility  in  hot  and  cold  water,  to  accom- 
plish their  separation,  and  carry  this  out.  Dissolve  some  of  the 
crystals  of  sodium  chromate  and  to  one  portion  of  the  solution 
add  silver  nitrate  solution,  then  acidify  with  nitric  acid.  To 
another  portion  add  sulfuric  acid  and  note  the  color  change, 
then  warm,  adding  alcohol  drop  by  drop,  and  again  note  the 
color  change.  Compare  these  colors  with  those  of  chromates, 
dichromates  and  any  chromium  salt,  and  account  for  each 
change  of  color  in  the  solution. 

107.  Chromyl  Chloride. 

Grind  together  2  grams  of  potassium  chromate  and  1  gram  of 
sodium  chloride,  place  a  small  portion  in  a  dry  test-tube  and  add 
two  or  three  drops  of  concentrated  sulfuric  acid  slowly.  Warm 
and  collect  a  drop  or  two  of  the  distillate  which  condenses  on 
the  walls  of  the  test-tube  by  pouring  it  from  the  test-tube  into 
a  beaker  containing  water.  Test  this  solution  with  litmus. 

107 


108  CHROMIUM 

Add  nitric  acid  and  silver  nitrate  solution  to  one  portion,  to 
another  portion  add  silver  nitrate  solution  without  acidifying. 
Explain  the  result  in  each  case. 

108.  The  Decomposition  of  Chromates  by  Heat. 

a.  Heat  a  small  portion  of  well  ground  potassium  dichromate 
in  a  dry,  hard  glass  test-tube.  Do  the  crystals  contain  water  of 
crystallization?  Is  the  substance  volatile  at  this  temperature? 
Heat  till  reaction  ceases  (how  may  that  point  be  determined?), 
testing  to  determine  what  gas  is  evolved.  Add  water  to  the 
residue  in  the  test-tube  after  cooling  and  heat  if  necessary  to 
accomplish  solution.  Filter.  What  is  the  precipitate?  Is 
it  soluble  in  sulfuric  acid?  What  is  the  nitrate? 

6.  To  5  c.c.  of  saturated  potassium  dichromate  solution  add 
concentrated  sulfuric  acid  slowly  till  a  permanent  precipitate 
forms.  What  is  the  precipitate?  Wet  a  piece  of  filter  paper 
in  this  mixture  and  explain  the  color  change.  Dry  the  precipi- 
tate by  decanting  as  much  of  the  liquid  as  possible  and  absorb- 
ing the  rest  on  a  porous  plate,  and  heat  some  of  this  dried 
material  in  a  dry  test-tube.  Is  it  volatile  ?  Does  it  decompose  ? 
Cool  and  treat  the  product  with  water  and  compare  with  the 
products  obtained  by  heating  potassium  dichromate. 

109.  Chrome  Alum. 

Calculate  the  amount  of  sulfuric  acid  necessary  to  make 
chrome  alum  from  5  grams  of  potassium  dichromate,  reducing 
with  alcohol.  Dissolve  5  grams  of  powdered  potassium  di- 
chromate in  a  small  amount  of  water,  add  the  calculated  amount 
of  sulfuric  acid  diluted  with  an  equal  volume  of  water,  heat  the 
solution  and  add  50  per  cent  alcohol,  drop  by  drop,  till  reduction 
is  complete.  Set  most  of  this  solution  aside  to  crystallize.  To 
a  small  portion  add  sodium  hydroxide  solution,  first  in  small 
quantities  then  in  excess,  and  explain  the  changes.  If  chrome 
iron  ore  were  fused  with  sodium  hydroxide  instead  of  the 
peroxide  (Exp.  106)  what  would  be  the  resulting  compound? 


CHAPTER  XXII 
MANGANESE 

110.  Compounds  Showing  the  Varying  Valence  of  Manganese. 

a.  Melt  4  grams  of  potassium  hydroxide  in  an  iron  crucible 
and  add  gradually  2  grams  of  powdered  manganese  dioxide, 
increasing  the  heat  as  necessary  to  keep  the  mixture  just  melted, 
and  stirring  with  an  iron  rod  or  file.  In  this  procedure  the 
manganese  undergoes  auto-oxidation,  part  being  oxidized  to- 
hexavalent  manganese  while  the  other  part  is  reduced  to- 
trivalent.  Write  the  equation,  assuming  that  the  oxides  are 
formed.  Which  of  these  oxides  would  be  acidic,  and  what  would 
be  the  effect  of  the  potassium  hydroxide  on  the  acid  oxide? 
Dissolve  the  fused  mixture  in  distilled  water  and  decant  through 
a  filter.  Keep  this  filtrate  for  d  and  the  insoluble  residue  for  b. 

6.  Wash  the  residue  with  distilled  water  until  the  wash  water 
is  colorless.  Some  of  the  original  manganese  dioxide  and  some 
manganic  oxide  will  be  present  in  this  residue.  Add  dilute 
sulfuric  acid.  Which  oxide  will  dissolve?  Allow  to  settle  and 
decant  the  clear  sulfuric  acid  solution  from  the  residue,  (save 
the  residue  for  c),  then  make  the  solution  alkaline  with  potas- 
sium hydroxide.  What  is  formed?  Keep  this  test  for  com- 
parison later. 

c.  Wash  the  residue  insoluble  in  sulfuric  acid  (what  is  it?) 
with  distilled  water  till  the  wash  water  is  colorless,  then  add 
dilute  sulfuric  acid,  warm  and  add  alcohol  or  sodium  sulfite 
solution,  drop  by  drop.  What  is  the  reaction?  Make  this  solu- 
tion alkaline  with  potassium  hydroxide  and  compare  with 
manganous  hydroxide,  made  by  adding  potassium  hydroxide  to 
manganous  sulfate,  and  with  the  manganic  hydroxide  prepared 
above.  Which  is  it?  Pour  off  as  much  of  the  supernatant 
liquid  as  possible  and  allow  the  precipitate  to  stand  several 

109 


110  MANGANESE 

hours,    and    compare   with    the   manganic    hydroxide    again. 
Account  for  the  change. 

d.  Pour  a  part  of  the  green  filtrate  from  the  fusion  mixture 
into  a  large  excess  of  water.     What  is  formed?     What  is  the 
oxidizing  agent?     The  reducing  agent?     Acidify  another  por- 
tion with  sulfuric  acid  and  account  for  the  change  in  color  of  the 
solution  and  for  the  precipitate  formed.     Add  sodium  sulfite 
solution  to  an  alkaline  and  an  acid  solution  of  potassium  per- 
manganate,  and  to  some  of  the  green  manganate  solution. 
Which  oxide  of  manganese  is  formed  when  alkaline  solutions 
of  manganates  and  permanganates  are  reduced?     When  an 
acid  solution  of  permanganate  (can  an  acid  solution  of  a  man- 
ganate be  formed?)  is  reduced  what  is  the  valence  of  the  man- 
ganese in  the  resulting  compound?     Devise  and  apply  a  test  to 
prove  your  answer  to  these  questions. 

e.  Add    dilute    nitric     acid    and    some    red    lead    to    a 
solution  of  mangarious  sulfate,  heat  to  boiling  then  allow  the 
precipitate  to  settle  and  note  and  account  for  the  color  of  the 
supernatant  liquid. 


CHAPTER  XXIII 
IRON 

111.  The  Preparation  of  Ammonium  Ferrous  Sulfate. 

Dissolve  5  grams  of  iron  nails  or  wire  by  heating  with  the 
calculated  amount  of  sulfuric  acid  diluted  with  four  times  its 
volume  of  water  (hood)  adding  water  if  necessary  to  replace  that 
lost  by  boiling,  but  keeping  the  final  volume  under  50  c.c. 
When  the  reaction  has  ceased  filter,  add  a  few  drops  of  dilute 
sulfuric  acid  and  the  calculated  amount  of  solid  ammonium 
sulfate  necessary  for  ferrous  ammonium  sulfate,  heat  to  boiling 
(add  water  only  if  necessary  to  dissolve  the  ammonium  sulfate 
at  the  boiling  temperature)  and  as  soon  as  solution  is  complete 
cool  rapidly  by  holding  under  running  water  and  stirring. 
Drain  the  crystals  from  the  supernatant  liquid,  wash  once  with 
a  small  amount  of  cold  distilled  water,  and  dry  the  crystals  on 
filter  paper  and  finally  by  standing  in  the  air.  Weigh  and 
calculate  the  per  cent,  yield  (i.e.  the  percentage  which  the 
amount  formed  is  of  the  theoretical  amount  obtainable  from  5 
grams  of  iron).  Dissolve  a  small  portion  and  test  the  solution 
with  litmus.  Save  this  preparation  for  Exps.  112  and  116. 

112.  Ammonium  Iron  Alum. 

a.  Dissolve  5  grams  of  the  ferrous  ammonium  sulfate  pre- 
pared in  Exp.  Ill  in  15  c.c.  of  distilled  water,  heat  to  boiling  and 
add  concentrated  nitric  acid  drop  by  drop  with  constant  boiling 
until  the  addition  of  more  nitric  acid  does  not  cause  the  forma- 
tion of  a  dark  brown  color  (to  what  is  this  color  due?).  Crys- 
tallize the  alum  by  cooling  and  evaporation,  and  remove  the 
crystals  from  the  liquid,  dry  on  filter  paper  and  place  in  a 
stoppered  bottle  for  use  in  Exp.  113.  Dissolve  a  small  portion 
and  test  the  solution  with  litmus.  Compare  with  the  ferrous 
salt. 

ill 


112  IRON 

b.  Test  solutions  of  ferrous  and  ferric  salts  with  ammonium 
hydroxide,  potassium  permanganate  solution  and  potassium 
iodide  solution.  Contrast  the  effects  of  the  two  iron  salts 
in  each  case.  Write  equations  for  each  raction. 

113.  Determination  of  the  Water  of  Crystallization  and  of  Ferric 
Oxide  in  Ammonium  Iron  Alum.     Quantitative. 

Weigh  a  clean  dry  porcelain  crucible,  place  in  it  0.2  to  0.5 
gram  of  the  ammonium  iron  alum  prepared  in  Exp.  112  (if 
these  crystals  have  become  opaque,  or  formed  a  brown  deposit 
on  the  surface  they  will  not  give  good  results  as  decomposition 
has  already  started)  and  determine  the  water  of  crystallization 
according  to  the  method  used  for  aluminium  alum,  Exp.  100. 
After  constant  weight  has  been  attained  at  the  low  temperature 
required  for  driving  off  the  water,  heat  the  crucible  slowly  to  red 
heat  with  the  direct  flame,  cool  and  weigh,  and  repeat  till  it 
comes  to  constant  weight.  The  residue  left  in  the  crucible  is 
ferric  oxide.  From  the  weights  of  the  water,  the  ferric  oxide 
and  the  original  weight  of  the  alum  taken  calculate  the  per  cent, 
of  water  and  iron  in  the  alum. 

114.  The   Hydrolysis    of    Ferric    Salts,    and  Colloidal  Ferric 
Hydroxide. 

a.  Heat  a  small  crystal  of  hydrated  ferric  chloride  in  a  dry, 
soft-glass  test-tube,  testing  the  gas  evolved  with  moist  litmus 
paper.     What  is  evolved?     Heat  till  the  test-tube  is  entirely 
dry,  cool,  add  water,  filter  and  test  the  filtrate  for  iron.     What  is 
the  insoluble  residue? 

b.  Make  50  c.c.  of  a  5  per  cent,  solution  of  ferric  chloride,  add 
sodium  carbonate  solution  as  long  as  the  precipitate  first  formed 
redissolves  (explain  the  phenomena  observed  here)  and  if  a 
permanent  precipitate  forms  add  dilute  hydrochloric  acid  till 
it  dissolves.     Dialyse  in  a  collodion  sack  (Exp.  71  6)  changing 
the  water  frequently  at  first  and  test  the  dialysate  for  chlorides. 
When  the  dialysate  shows  only  traces  of  chlorides  after  standing 
in  contact  with  the  sack  for  about  an  hour  remove  the  sack  and 


FERROUS  AMMONIUM  SULFATE  113 

test  the  solution  of  colloidal  iron  for  chlorides.  Add  a  concen- 
trated solution  of  sodium  chloride  to  this  solution.  Is  this  a 
physical  or  a  chemical  change?  Add  a  dilute  acid  to  another 
portion  of  the  colloidal  iron  hydroxide  solution  and  account  for 
the  change  in  color. 

c.  Recall  the  effect  of  ferrous  and  ferric  sulfate  solutions  on 
litmus.  In  which  valence  does  iron  show  most  marked  metal- 
lic properties? 

115.  Cast  Iron  and  Steel. 

Dissolve  small  samples  of  cast  iron  and  steel  in  dilute  hydro- 
chloric acid,  testing  the  gases  evolved  for  hydrogen  sulfide. 
Note  the  amount  and  appearance  of  the  residue  in  each  case. 
Filter  and  mix  the  dried  residues  with  copper  oxide  and  heat 
each  separately  in  a  hard  glass  test-tube,  testing  the  evolved 
gas  for  carbon  dioxide.  Dissolve  about  a  gram  of  steel  filings 
by  stirring  with  potassium  cupric  chloride  solution  (what  is  the 
white  precipitate  that  forms  and  redissolves?).  Note  the 
appearance  of  the  residue  and  test  the  gases  evolved  when  it  is 
heated  with  copper  oxide.  What  elements  besides  iron  are 
present  in  cast  iron  and  steel? 

116.  The  Oxidation  of  Ferrous  Ammonium  Sulfate  by  Potas- 
sium Permanganate  and  by  Potassium  Dichromate.    Quan- 
titative. 

Make  0.1  formular  solutions  of  potassium  permanganate  and 
potassium  dichromate  by  weighing  accurately  the  amount  re- 
quired for  100  c.c.,  placing  this  amount  in  the  measuring 
cylinder,  dissolving  in  distilled  water,  then  adding  water  to 
make  just  100  c.c.  (be  careful  that  none  of  the  solution  is  lost). 
The  rest  of  this  experiment  must  be  performed  in  one  laboratory 
period.  Make  100  c.c.  of  a  0.1  formular  ammonium  ferrous 
sulfate  solution,  weighing  out  the  salt  prepared  in  Exp.  Ill 
accurately  (remember  that  these  crystals  contain  water)  and 
place  exactly  50  c.c.  of  this  solution  in  each  of  two  beakers. 
Add  to  each  solution  5  c.c.  of  dilute  sulfuric  acid,  heat  to  boiling 
and  to  the  first  beaker  add  the  prepared  solution  of  potassium 


114  IRON 

permanganate  from  a  burette  until  the  purple  color  of  the  per- 
manganate just  becomes  permanent,  and  record  the  amount  of 
permanganate  solution  required.  To  the  second  beaker  con- 
taining ammonium  ferrous  sulfate  heated  to  boiling  add  potas- 
sium dichromate  solution  from  a  burette;  1  c.c.  at  a  time, 
testing  a  drop  of  the  solution  after  each  addition  by  bringing 
it  in  contact  with  a  drop  of  potassium  ferricyanide  solution  on  a 
white  dish  or  test  plate.  As  long  as  a  blue  precipitate  or  solu- 
tion is  formed  on  mixing  these  drops  there  is  still  some  unchanged 
ferrous  sulfate  in  the  solution,  and  potassium  dichromate 
solution  should  be  added  only  until  the  test  drop  no  longer  gives 
a  blue  color.  Record  the  amount  of  potassium  dichromate 
solution  required.  How  do  these  two  oxidizing  agents  compare 
in  oxidizing  power,  or  in  the  amount  of  oxidation  accomplished 
by  equal  molecular  quantities  of  each  of  the  salts?  In  order  to 
make  a  normal  solution  of  each  so  that  the  oxygen  available 
for  oxidations  should  be  equivalent  to  1  gram  of  hydrogen  per 
liter,  how  much  of  each  salt  would  be  required  for  a  liter  of 
solution? 


INDEX 


Acetylene,  79 

Air,  determination  of  oxygen  in,  68; 
dissolved  in  water,  composition 
of,  69 

Alcohol,  the  oxidation  products  of, 
79 

Alum,  102;  ammonium  iron,   111 

Aluminates,  102 

Aluminium,  hydroxide,  102;  in  Kao- 
lin, 103 

Ammonia,  63 ;  weight  of  a  liter  of,  64 

Ammonium, ferrous  sulfate,  111;  hy- 
droxide, 64;  iron  alum,  111 

Antimony,  75 

Arsenic,  73 

Balance,  14 

Barium   bicarbonate,    solubility   of, 

96;  hydroxide,   solubility  of,   95; 

sulfate,  solubility  of,  96 
Bismuth,  75 

Boiling  point  as  a  test  for  purity,  16 
Borates  of  the  heavy  metals,  84 
Boric  acid,  preparation  from  borax, 

83;  anhydride,  83 
Boron,  preparation,  83 
Bromine,  preparation  of,  49 
Bunsen  burner,  10 
Burette,  11 

Calcium,  bicarbonate,  solubility  of, 
96;  carbonate,  95;  chloride,  97; 
hydroxide,  solubility  of,  95;  sul- 
fate, solubility  of,  96 

Carbon,  76;  dioxide,  76;  dioxide, 
density  of,  77;  monoxide,  77 

Catalysis,  21 

Chlorine,  density  of,  38;  preparation 
of,  38;  properties  of,  41 

Chromate,  sodium,  107 

Chromates,  decomposition  of,  108 

Chrome,  alum,  108;  iron  ore,  107 

Chromyl  chloride,  107 

Colloidal,  ferric  hydroxide,  112;  sil- 
icic acid,  81 ;  sulfur,  56 

Combining  weight,  of  copper,  29; 
of  tin,  67 

Concentration,  of  a  solution  of  acid 
or  base,  59;  on  the  speed  of  a  re- 
action, effect  of,  23 


Copper,  92;  combining  weight  of,  29 
Cupric  compounds,  92 
Cuprous  compounds,  92 


Definite  proportion,  law  of,  20 
Diffusion  of  gases,  rate  of,  28 
Dulong  and  Petit 's  law,  106 


Electromotive  series,  93 
Elements  and  compounds,  18 
Equilibrium,  53 
Ethylene,  79 

Ferric,    hydroxide,     colloidal,     112; 

salts,  hydrolysis  of,  112 
Ferrous    ammonium    sulfate,     111; 

oxidation  of,  113 

General  directions,  1 

Hydrochloric  acid,  42 

Hydrofluoric  acid,  52 

Hydrogen,  preparation  and  proper- 
ties, 26 

Hydrogen  peroxide,  concentration 
of,  35;  preparation  and  properties 
of,  35 

Hydrogen  sulfide,  56 

Hard  water,  77 

Hydrolysis,  85 

Iodine,  52 

Indicators,  86;  quantitative,  88 
lonization,  44;  85 

Iron  alum,  111;  water  of  crystalli- 
zation and  ferric  oxide  in,  112 
Iron,  cast,  113 

Law  of  definite  proportion,  20 
Law  of  multiple  proportion,  36 
Lead,  properties  of,  104;  oxides  of, 

104;  salts,  properties  of,  104 
Lime,  95 

Magnesium,   equivalent  weight  of, 

98;  carbonate,  99;  oxide,  99 
Manganese,  varying  valence  of,  109 
Measuring  instruments,  11 
Melting  point  a  test  for  purity,  17 
Mercuric  nitrate,  preparation  of,  101 


115 


116 


INDEX 


Mercurous  nitrate,  preparation  of, 

100 

Methane,  78 
Multiple  proportion,  law  of,  36 

Neutralization,  59 
Nitrates,  behavior  on  heating,  66 
Nitric  acid,  65;  oxidation  with,  67 
Nitric  oxide,  65 

Nitrites,  behavior  on  heating,  66 
Nitrogen,  preparation  and  proper- 
ties, 63 
Nitrous  oxide,  66 

Orthophosphates,  the  preparation 
of  primary,  secondary  and  terti- 
ary, 72 

Oxygen,  determination  of,  in  the  air, 
68;  preparation  and  properties  of, 
21;  weight  of  a  liter  of,  23 

Phosphates  of  sodium,  72 

Phosphorus,  oxides  of,  71 

Pipette,  11 

Potassium,  acid  sulf ate,  62;  bromate, 
50;  chlorate,  47;  dichromate,  oxi- 
dation by,  113;  dichromate,  sol- 
ubility of,  32;  hypochlorite,  46; 
nitrate,  90;  per  chlorate,  prepara- 
tion of,  47;  perchlorate,  purity 
of,  48;  permanganate,  oxidation 
by,  113 

Primary  orthophosphates,  72 

Pure  substances  and  mixtures,  16 

Quantitative  experiments,  directions 
for,  3 

Rock  salt,  purification  of,  45 

Secondary  orthophosphates,  72 

Silicic  acid,  81 

Silicon  tetrafluoride,  81 


Silver,  chloride  emulsion,  94;  oxide, 
94 

Soap,  78 

Sodium,  carbonate  by  the  Solvay 
process,  89;  chloride,  purification 
of,  45;  chromate  from  chrome  iron 
ore,  107;  sulf  ate,  89;  thiosulfate, 
preparation  and  properties  of,  58 

Solubility  of  salts  in  water,  quant.,  32 

Solutions,  supersaturated,  33 

Standard  conditions  for  gases,  5 

Stannic  tin,  105 

Stannous  tin,  105 

Steel,  113 

Strontium,  bicarbonate,  solubility 
of,  96;  hydroxide,  solubility  of, 
95;  sulf  ate,  solubility  of,  96 

Sulfur,  chemical  properties  of,  56; 
physical  properties  of,  55;  dioxide, 
57 

Sulfuric  acid,  59 

Table,  of  barometric  corrections,  6;  of 
densities  and  composition  of  water 
solutions  of  acids  and  ammonia,  7; 
of  solubility  and  density  of  gases, 
9;  of  solubility  of  salts,  8;  of  vapor 
pressures,  6 

Tertiary  orthophosphates,  72 
Tin,   atomic  weight  of,    105;   com- 
bining weight  of,  67,  105;  stannous 
and  stannic,  105 

Water,  composition  of,  31;  hard,  77; 
of  crystallization  in  alum,  102;  of 
crystallization  in  iron  alum,  112; 
of  hydration,  34 

Waters,  natural,  34 

Zinc,  equivalent  weight  of,  99;  prop- 
erties of,  99;  sulf  ate,  preparation 
of,  100 


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